AP Chem Notes chapters 7 8 9

ReadDo Problems for Chapters 7 8 9Chapter 7 Atomic Structure and Periodicity

Electromagnetic Radiation Radiant energy that exhibits wavelength-like behavior and travels

through space at the speed of light in a vacuum

Label the parts of the graph below wavelength amplitude node

Waves have 3 primary characteristics 1 Wavelength (λ) distance between two peaks in a wave2 Frequency (ν) number of waves per second that pass a given point in space3

Speed (c) speed of light is 29979 x 108 meterssecond (ms)

Wavelength and frequency can be inter-converted

c = λν

ν = frequency (s-1 1s Hz cycs or wavess ) 1 hertz (Hz) = 1s

λ = wavelength (m) c = speed of light (ms)

Max Planck (1858-1947)

An object can gain or lose energy by absorbing or emitting radiant energy in QUANTA

Transfer of energy is quantized and can only occur in discrete units called quanta

∆ E=hν=hcλ

ΔE = change in energy (J)h = Planckrsquos constant 6626 x 10-34 J s ν = frequency (s-1) λ = wavelength (m)

Photoelectric EffectAlbert Einstein (1879-1955)

Experiment demonstrates the particle nature of light Einstein called light particles ldquophotonsrdquo

Classical theory said that E of ejected electron should increase with increase in light intensityhellip but

No e- observed until light of a certain minimum E is used Number of e- ejected depends on light intensity

Energy has mass E = mc2

E = energy m = mass c = speed of light

Classical physics viewed light as a form of energy so it predicted that a metal would eventually collect enough energy to eject an electron regardless of the wavelength of the light

The emission of electrons from a metal when light shines on the metal

For a given metal no electrons were emitted if the frequency of light was below a ldquothresholdrdquo value regardless of how long the light was shining on the metal

Ephoton = hcλ

mphoton = hλc

Hence the dual nature of light or wave-particle duality

Line Spectra of excited atomsExcited atoms emit light of only certain wavelengthsThe wavelengths of emitted light depend on the element Observe spectrum tubes with diffraction gratings (spectroscopes)

Visible lines in H atom spectrum are called the BALMER series

Continuous spectrum Contains all the wavelengths of light

Line (discrete) spectrum Contains only some of the wavelengths of light

Niels Bohr (1885-1962) Bohrrsquos greatest contribution to science was in building a simple model of the atom It was based on an understanding of the SHARP LINE SPECTRA of excited atoms

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

+Electronorbit

1 Any orbit should be possible and so is any energy2 But a charged particle moving in an electric field should emit energy

End result should be destruction

Bohr said classical view is wrong Need a new theory mdash now called QUANTUM or WAVE MECHANICS e- can only exist in certain discrete orbits mdash called stationary states e- is restricted to QUANTIZED energy states Energy of state = - Cn2 where C is a constant amp where n = quantum no = 1 2 3 4

Bohr Model The e- in a hydrogen atom moves around the nucleus only in certain allowed circular

orbits

E = energy of the levels in the H-atomz = nuclear charge (for H z = 1)n = an integer

E becomes more negative as the electron moves closer to the nucleus

Ground State The lowest possible energy state for an atom (n = 1)

Shielding Effect Higher energy electrons are in outer shells and are easier to remove because they are more shielded from the positive nucleus by inner shell electrons

Energy Changes in the Hydrogen AtomΔE = Efinal state ndash Einitial state

λ= hcΔE

Failure of Bohr ModelFailure of Bohr Model The Bohr Model of the atom paved the way for the Quantum Mechanical Theory

but current theory is in no way derived from the Bohr Model of the atom The Bohr Model of the Atom was fundamentally incorrect--electrons do not move in

circular orbits about the nucleus

Umhellipso what is quantum mechanics

Louis deBroglie (1892-1987)

de Broglie (1924) proposed that all moving objects have wave properties

λ = hm v

λ = wavelength (m)h = Planckrsquos constant 6626 x 10-34 J s = kg m2 s-1 m = mass (kg)v = velocity (ms) ndash do not confuse with nu (frequency) or c (speed of light) ndash do not confuse with nu (frequency) or c (speed of light)

See Sample Exercise 73 on page 298

Baseball (0100 kg) at 35 ms

λ = 19 x 10-34 m

e- with velocity = 10 x 107 ms

λ = 727 x 10-11 m

Erwin Schrodinger 1887-1961

Transfer of energy is quantized and can only occur in discrete units called quanta

∆ E=hν=hcλ

ΔE = change in energy (J)h = Planckrsquos constant 6626 x 10-34 J s ν = frequency (s-1) λ = wavelength (m)

Photoelectric EffectAlbert Einstein (1879-1955)

Experiment demonstrates the particle nature of light Einstein called light particles ldquophotonsrdquo

Classical theory said that E of ejected electron should increase with increase in light intensityhellip but

No e- observed until light of a certain minimum E is used Number of e- ejected depends on light intensity

Energy has mass E = mc2

E = energy m = mass c = speed of light

Classical physics viewed light as a form of energy so it predicted that a metal would eventually collect enough energy to eject an electron regardless of the wavelength of the light

The emission of electrons from a metal when light shines on the metal

For a given metal no electrons were emitted if the frequency of light was below a ldquothresholdrdquo value regardless of how long the light was shining on the metal

Ephoton = hcλ

mphoton = hλc

Hence the dual nature of light or wave-particle duality

Line Spectra of excited atomsExcited atoms emit light of only certain wavelengthsThe wavelengths of emitted light depend on the element Observe spectrum tubes with diffraction gratings (spectroscopes)

Visible lines in H atom spectrum are called the BALMER series

Continuous spectrum Contains all the wavelengths of light

Line (discrete) spectrum Contains only some of the wavelengths of light

Niels Bohr (1885-1962) Bohrrsquos greatest contribution to science was in building a simple model of the atom It was based on an understanding of the SHARP LINE SPECTRA of excited atoms

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

+Electronorbit

1 Any orbit should be possible and so is any energy2 But a charged particle moving in an electric field should emit energy

End result should be destruction

Bohr said classical view is wrong Need a new theory mdash now called QUANTUM or WAVE MECHANICS e- can only exist in certain discrete orbits mdash called stationary states e- is restricted to QUANTIZED energy states Energy of state = - Cn2 where C is a constant amp where n = quantum no = 1 2 3 4

Bohr Model The e- in a hydrogen atom moves around the nucleus only in certain allowed circular

orbits

E = energy of the levels in the H-atomz = nuclear charge (for H z = 1)n = an integer

E becomes more negative as the electron moves closer to the nucleus

Ground State The lowest possible energy state for an atom (n = 1)

Shielding Effect Higher energy electrons are in outer shells and are easier to remove because they are more shielded from the positive nucleus by inner shell electrons

Energy Changes in the Hydrogen AtomΔE = Efinal state ndash Einitial state

λ= hcΔE

Failure of Bohr ModelFailure of Bohr Model The Bohr Model of the atom paved the way for the Quantum Mechanical Theory

but current theory is in no way derived from the Bohr Model of the atom The Bohr Model of the Atom was fundamentally incorrect--electrons do not move in

circular orbits about the nucleus

Umhellipso what is quantum mechanics

Louis deBroglie (1892-1987)

de Broglie (1924) proposed that all moving objects have wave properties

λ = hm v

λ = wavelength (m)h = Planckrsquos constant 6626 x 10-34 J s = kg m2 s-1 m = mass (kg)v = velocity (ms) ndash do not confuse with nu (frequency) or c (speed of light) ndash do not confuse with nu (frequency) or c (speed of light)

See Sample Exercise 73 on page 298

Baseball (0100 kg) at 35 ms

λ = 19 x 10-34 m

e- with velocity = 10 x 107 ms

λ = 727 x 10-11 m

Erwin Schrodinger 1887-1961

Ephoton = hcλ

mphoton = hλc

Hence the dual nature of light or wave-particle duality

Line Spectra of excited atomsExcited atoms emit light of only certain wavelengthsThe wavelengths of emitted light depend on the element Observe spectrum tubes with diffraction gratings (spectroscopes)

Visible lines in H atom spectrum are called the BALMER series

Continuous spectrum Contains all the wavelengths of light

Line (discrete) spectrum Contains only some of the wavelengths of light

Niels Bohr (1885-1962) Bohrrsquos greatest contribution to science was in building a simple model of the atom It was based on an understanding of the SHARP LINE SPECTRA of excited atoms

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

+Electronorbit

1 Any orbit should be possible and so is any energy2 But a charged particle moving in an electric field should emit energy

End result should be destruction

Bohr said classical view is wrong Need a new theory mdash now called QUANTUM or WAVE MECHANICS e- can only exist in certain discrete orbits mdash called stationary states e- is restricted to QUANTIZED energy states Energy of state = - Cn2 where C is a constant amp where n = quantum no = 1 2 3 4

Bohr Model The e- in a hydrogen atom moves around the nucleus only in certain allowed circular

orbits

E = energy of the levels in the H-atomz = nuclear charge (for H z = 1)n = an integer

E becomes more negative as the electron moves closer to the nucleus

Ground State The lowest possible energy state for an atom (n = 1)

Shielding Effect Higher energy electrons are in outer shells and are easier to remove because they are more shielded from the positive nucleus by inner shell electrons

Energy Changes in the Hydrogen AtomΔE = Efinal state ndash Einitial state

λ= hcΔE

Failure of Bohr ModelFailure of Bohr Model The Bohr Model of the atom paved the way for the Quantum Mechanical Theory

but current theory is in no way derived from the Bohr Model of the atom The Bohr Model of the Atom was fundamentally incorrect--electrons do not move in

circular orbits about the nucleus

Umhellipso what is quantum mechanics

Louis deBroglie (1892-1987)

de Broglie (1924) proposed that all moving objects have wave properties

λ = hm v

λ = wavelength (m)h = Planckrsquos constant 6626 x 10-34 J s = kg m2 s-1 m = mass (kg)v = velocity (ms) ndash do not confuse with nu (frequency) or c (speed of light) ndash do not confuse with nu (frequency) or c (speed of light)

See Sample Exercise 73 on page 298

Baseball (0100 kg) at 35 ms

λ = 19 x 10-34 m

e- with velocity = 10 x 107 ms

λ = 727 x 10-11 m

Erwin Schrodinger 1887-1961

Niels Bohr (1885-1962) Bohrrsquos greatest contribution to science was in building a simple model of the atom It was based on an understanding of the SHARP LINE SPECTRA of excited atoms

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

+Electronorbit

1 Any orbit should be possible and so is any energy2 But a charged particle moving in an electric field should emit energy

End result should be destruction

Bohr said classical view is wrong Need a new theory mdash now called QUANTUM or WAVE MECHANICS e- can only exist in certain discrete orbits mdash called stationary states e- is restricted to QUANTIZED energy states Energy of state = - Cn2 where C is a constant amp where n = quantum no = 1 2 3 4

Bohr Model The e- in a hydrogen atom moves around the nucleus only in certain allowed circular

orbits

E = energy of the levels in the H-atomz = nuclear charge (for H z = 1)n = an integer

E becomes more negative as the electron moves closer to the nucleus

Ground State The lowest possible energy state for an atom (n = 1)

Shielding Effect Higher energy electrons are in outer shells and are easier to remove because they are more shielded from the positive nucleus by inner shell electrons

Energy Changes in the Hydrogen AtomΔE = Efinal state ndash Einitial state

λ= hcΔE

Failure of Bohr ModelFailure of Bohr Model The Bohr Model of the atom paved the way for the Quantum Mechanical Theory

but current theory is in no way derived from the Bohr Model of the atom The Bohr Model of the Atom was fundamentally incorrect--electrons do not move in

circular orbits about the nucleus

Umhellipso what is quantum mechanics

Louis deBroglie (1892-1987)

de Broglie (1924) proposed that all moving objects have wave properties

λ = hm v

λ = wavelength (m)h = Planckrsquos constant 6626 x 10-34 J s = kg m2 s-1 m = mass (kg)v = velocity (ms) ndash do not confuse with nu (frequency) or c (speed of light) ndash do not confuse with nu (frequency) or c (speed of light)

See Sample Exercise 73 on page 298

Baseball (0100 kg) at 35 ms

λ = 19 x 10-34 m

e- with velocity = 10 x 107 ms

λ = 727 x 10-11 m

Erwin Schrodinger 1887-1961

Umhellipso what is quantum mechanics

Louis deBroglie (1892-1987)

de Broglie (1924) proposed that all moving objects have wave properties

λ = hm v

λ = wavelength (m)h = Planckrsquos constant 6626 x 10-34 J s = kg m2 s-1 m = mass (kg)v = velocity (ms) ndash do not confuse with nu (frequency) or c (speed of light) ndash do not confuse with nu (frequency) or c (speed of light)

See Sample Exercise 73 on page 298

Baseball (0100 kg) at 35 ms

λ = 19 x 10-34 m

e- with velocity = 10 x 107 ms

λ = 727 x 10-11 m

Erwin Schrodinger 1887-1961

Schrodinger applied idea of e- behaving as a wave to the problem of electrons in atomsWave Equation Based on the wave properties of the atom

Ĥ ψ = E ψψ = wave function Ĥ = mathematical operatorE = total energy of the atomA specific wave function is often called an orbital Electrons move in orbitals not Bohr orbits

Werner Heisenberg 1901-1976

Heisenberg Uncertainty Principle Problem of defining nature of electrons in atoms solved by Werner Heisenberg Cannot simultaneously define the position and momentum (p = mbullv) of an electron We define e- energy exactly but accept the limitation that we do not know exact

position

x = positionmv = momentum (p)h = Planckrsquos constant

The more accurately we know a particlersquos position the less accurately we can know its momentum

Probability Distribution square of the wave function (ψ2) probability of finding an electron at a given position Radial probability distribution is the probability distribution in each spherical shell

Quantum Numbers describe the location of the eQuantum Numbers describe the location of the e--

symbol

name possible values what it tells about electron

cloud

n principal quantum number

1-7 size(energy level)

langular momentum(sublevel) quantum

numberl =0 to n-1 shape

l = 0 (s sublevel) l = 1 (p sublevel) l = 2 (d sublevel) l = 3 (f sublevel)

ml magnetic quantum number m

l = l to -l 3D orientation

in space

ms spin quantum number plusmn 12

spin states of electron

Pauli Exclusion Principle In a given atom no two electrons can have the same set of four quantum numbers

(n l ml ms) Therefore an orbital can hold only two electrons and they must have opposite spins

Shapes of Electron Orbitals

The following are graphical representations of the shapes of the orbitals within atoms Remember these shapes do not exist but are mainly three-dimensional probabilities of locating the electron Also remember a maximum of two electrons can occupy any of the following shapes

How many different kinds of orbitals does each energy level have n =

the number of the energy level

n2 = the number of orbitals in an energy level 2n2 = the number of electrons in an energy level

s-orbital (spherical)The following picture shows the relative sizes of 1s 2s and 3s orbitals

p-orbital (dumbbell)

3 orbitals px py pz - The three p orbitals lie 90deg apart in space

When l = 1 there is a PLANAR NODE thru the nucleus When n = 2 then l = 0 and 1Therefore in n = 2 shell there are 2 types of orbitals mdash 2 subshells For l = 0 m

l = 0 this is a s subshell

For l = 1 ml = -1 0 +1 this is a p subshell with 3 orbitals

d-orbital (double dumbbell and dumbbelldonut)5 orbitalsWhen n = 3 what are the values of l

l = 0 1 2 and so there are 3 subshells in the shell

For l = 0 ml = 0 ---gt s subshell with single orbital

For l = 1 ml = -1 0 +1 ---gt p subshell with 3 orbitals

For l = 2 ml = -2 -1 0 +1 +2---gt d subshell with 5 orbitals

s orbitals have no planar node (l = 0) and so are spherical

p orbitals have l = 1 and have 1 planar node and so are ldquodumbbellrdquo shaped

This means d orbitals (with l = 2) have 2 planar nodes

f-orbital (quadruple dumbbell and dumbbelldouble donut)7 orbitalsWhen n = 4 When n = 4 l = 0 1 2 3 so there are 4 subshells in the shell = 0 1 2 3 so there are 4 subshells in the shell

For For l = 0 = 0 ml = 0 ---gt s subshell with single orbital = 0 ---gt s subshell with single orbital

For For l = 1 = 1 ml = -1 0 +1 ---gt p subshell with 3 orbitals = -1 0 +1 ---gt p subshell with 3 orbitals

For For l = 2 = 2 ml = -2 -1 0 +1 +2 = -2 -1 0 +1 +2 ---gt d subshell with 5 orbitals ---gt d subshell with 5 orbitals

For For l = 3 = 3 ml = -3 -2 -1 0 +1 +2 +3 = -3 -2 -1 0 +1 +2 +3 ---gt f subshell with 7 orbitals ---gt f subshell with 7 orbitals

Aufbau principleAs protons are added one by one to the nucleus to build up the elements electrons

are similarly added to these hydrogen-like orbitals

electron filling order (aufbau)

Hundrsquos Rule The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate (same energy) orbitals

Valence electrons electrons in the outermost principle quantum level of an atomThe inner electron are called core electrons

atom valence electronsCa 2N 5Br 7Arrangement of Electrons in Atoms

Shells (energy levels) = ndarr

Sublevels = ldarr

Orbitals = ml

Assigning electrons to atomsAssigning electrons to atomsElectrons generally assigned to orbitals of successively higher energyFor H atoms - Cn2 E depends only on nFor many-electron atoms energy depends on both n and l

Effective Nuclear Charge Zeff

Zeff is the nuclear charge experienced by the outermost electrons Explains why E(2s) lt E(2p)

Zeff increases across a period owing to incomplete shielding by inner electrons

Estimate Zeff by --gt [ Z - (no inner electrons) ]

Charge felt by 2s e- in Li Zeff = 3 - 2 = 1Be Zeff = 4 - 2 = 2B Zeff = 5 - 2 = 3and so on

Writing Atomic Electron Configurations

Orbital Notation Each electron is represented by a half-arrow

When electrons share an orbital opposite spins are represented by half-arrows going in opposite directions

Examples Write the orbital notation for the following atoms

N (7 e-) ____ ____ ____ ____ ____

1s 2s 2p

P ( ___ e-)

Electron Configuration Notation Write the energy level and orbital (in increasing energy) Write the number of electrons in the sublevel with a superscript

Examples Write the electron configuration notation for the following atoms

N

P

Tc

Electron Configuration Notation (Noble Gas Shorthand) Start with the noble gas ldquocloset without going overrdquo in brackets Continue normal electron configuration

Examples Write the electron configuration notation for the PREVIOUS atoms (use the noble gas shorthand)

Lewis Dot Diagrams

Valence electrons (e- in the outermost energy level) are the only electrons that are lost or gained in ordinary chemical reactions

Lewis Dot Diagrams only include electrons in valence (highest energy level) Write the letter and dots around the outside (2 on each side maximum of 8)

Examples Draw the Lewis Dot Diagram for

N Ca Cl

Ion configurations To form cations from elements remove 1 or more e- from subshell of highest n [or highest (n + l)]

P [Ne] 3s2 3p3 - 3e- ---gt P3+ [Ne] 3s2 3p0

For transition metals remove ns electrons and then (n - 1) electronsFe [Ar] 4s2 3d5

loses 2 electrons ---gt Fe2+ [Ar] 4s0 3d6

loses 3 electrons ---gt Fe3+

[Ar] 4s0 3d5

Information Contained on the Periodic Table1 Each group member has the same valence electron configuration

(these electrons primarily determine an atomrsquos chemistry)

2 The electron configuration of any representative element (see order)

3 Certain groups have special names (alkali metals halogens etc)02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

4 Metals and nonmetals are characterized by their chemical and physical properties

Periodic Trends overall change in physical property through row or group on the periodic

table These are general trends (there are exceptions)

Trend 1 Atomic Radius

Trend 2 Ionic Radius

Cations (positive ions) generally are smaller than their respective atoms

Anions (negative ions) generally are larger than their respective atoms

Trend 3 ndash Ionization Energy (IE)

Trend 4 ndash Electronegativity (en)

Redox Reactions Why do metals lose electrons in their reactions

Why does Mg form Mg2+ ions and not Mg3+Why do nonmetals take on electrons

Electron Affinity A few elements GAIN electrons to form anions

Electron affinity is the energy involved when an anion loses an electron A -(g) rarr A(g) + e- EA = ΔE

EA = 141 kJΔE is ENDOthermic because O has an affinity for an e-

O [He] - ion

[He] O atom

EA = 0 kJΔE is zero for N-

due to electron-electron repulsions

Trend 5 Electron AffinityAffinity for electron increases across a period (EA becomes more negative)

Reason

Affinity decreases down a group (EA becomes less negative)

Reason

[He] N- ion

[He] N atom

Atom EAF +328 kJCl+349 kJBr+325 kJI +295 kJ

1 2 3 4 5 6 7S1

S2

S3

S4

0

50

100

150

200

250

300

350

Elec

tron

affi

nity

(kJm

ol)

Group

Period

02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

Increasing Periodic Trends

electronegativity ionization energy ionic radii electron affinityatomic radii

ionic amp atomic radii

ionization energyelectron affinity amp electronegativity

p-orbital (dumbbell)

3 orbitals px py pz - The three p orbitals lie 90deg apart in space

When l = 1 there is a PLANAR NODE thru the nucleus When n = 2 then l = 0 and 1Therefore in n = 2 shell there are 2 types of orbitals mdash 2 subshells For l = 0 m

l = 0 this is a s subshell

For l = 1 ml = -1 0 +1 this is a p subshell with 3 orbitals

d-orbital (double dumbbell and dumbbelldonut)5 orbitalsWhen n = 3 what are the values of l

l = 0 1 2 and so there are 3 subshells in the shell

For l = 0 ml = 0 ---gt s subshell with single orbital

For l = 1 ml = -1 0 +1 ---gt p subshell with 3 orbitals

For l = 2 ml = -2 -1 0 +1 +2---gt d subshell with 5 orbitals

s orbitals have no planar node (l = 0) and so are spherical

p orbitals have l = 1 and have 1 planar node and so are ldquodumbbellrdquo shaped

This means d orbitals (with l = 2) have 2 planar nodes

f-orbital (quadruple dumbbell and dumbbelldouble donut)7 orbitalsWhen n = 4 When n = 4 l = 0 1 2 3 so there are 4 subshells in the shell = 0 1 2 3 so there are 4 subshells in the shell

For For l = 0 = 0 ml = 0 ---gt s subshell with single orbital = 0 ---gt s subshell with single orbital

For For l = 1 = 1 ml = -1 0 +1 ---gt p subshell with 3 orbitals = -1 0 +1 ---gt p subshell with 3 orbitals

For For l = 2 = 2 ml = -2 -1 0 +1 +2 = -2 -1 0 +1 +2 ---gt d subshell with 5 orbitals ---gt d subshell with 5 orbitals

For For l = 3 = 3 ml = -3 -2 -1 0 +1 +2 +3 = -3 -2 -1 0 +1 +2 +3 ---gt f subshell with 7 orbitals ---gt f subshell with 7 orbitals

Aufbau principleAs protons are added one by one to the nucleus to build up the elements electrons

are similarly added to these hydrogen-like orbitals

electron filling order (aufbau)

Hundrsquos Rule The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate (same energy) orbitals

Valence electrons electrons in the outermost principle quantum level of an atomThe inner electron are called core electrons

atom valence electronsCa 2N 5Br 7Arrangement of Electrons in Atoms

Shells (energy levels) = ndarr

Sublevels = ldarr

Orbitals = ml

Assigning electrons to atomsAssigning electrons to atomsElectrons generally assigned to orbitals of successively higher energyFor H atoms - Cn2 E depends only on nFor many-electron atoms energy depends on both n and l

Effective Nuclear Charge Zeff

Zeff is the nuclear charge experienced by the outermost electrons Explains why E(2s) lt E(2p)

Zeff increases across a period owing to incomplete shielding by inner electrons

Estimate Zeff by --gt [ Z - (no inner electrons) ]

Charge felt by 2s e- in Li Zeff = 3 - 2 = 1Be Zeff = 4 - 2 = 2B Zeff = 5 - 2 = 3and so on

Writing Atomic Electron Configurations

Orbital Notation Each electron is represented by a half-arrow

When electrons share an orbital opposite spins are represented by half-arrows going in opposite directions

Examples Write the orbital notation for the following atoms

N (7 e-) ____ ____ ____ ____ ____

1s 2s 2p

P ( ___ e-)

Electron Configuration Notation Write the energy level and orbital (in increasing energy) Write the number of electrons in the sublevel with a superscript

Examples Write the electron configuration notation for the following atoms

N

P

Tc

Electron Configuration Notation (Noble Gas Shorthand) Start with the noble gas ldquocloset without going overrdquo in brackets Continue normal electron configuration

Examples Write the electron configuration notation for the PREVIOUS atoms (use the noble gas shorthand)

Lewis Dot Diagrams

Valence electrons (e- in the outermost energy level) are the only electrons that are lost or gained in ordinary chemical reactions

Lewis Dot Diagrams only include electrons in valence (highest energy level) Write the letter and dots around the outside (2 on each side maximum of 8)

Examples Draw the Lewis Dot Diagram for

N Ca Cl

Ion configurations To form cations from elements remove 1 or more e- from subshell of highest n [or highest (n + l)]

P [Ne] 3s2 3p3 - 3e- ---gt P3+ [Ne] 3s2 3p0

For transition metals remove ns electrons and then (n - 1) electronsFe [Ar] 4s2 3d5

loses 2 electrons ---gt Fe2+ [Ar] 4s0 3d6

loses 3 electrons ---gt Fe3+

[Ar] 4s0 3d5

Information Contained on the Periodic Table1 Each group member has the same valence electron configuration

(these electrons primarily determine an atomrsquos chemistry)

2 The electron configuration of any representative element (see order)

3 Certain groups have special names (alkali metals halogens etc)02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

4 Metals and nonmetals are characterized by their chemical and physical properties

Periodic Trends overall change in physical property through row or group on the periodic

table These are general trends (there are exceptions)

Trend 1 Atomic Radius

Trend 2 Ionic Radius

Cations (positive ions) generally are smaller than their respective atoms

Anions (negative ions) generally are larger than their respective atoms

Trend 3 ndash Ionization Energy (IE)

Trend 4 ndash Electronegativity (en)

Redox Reactions Why do metals lose electrons in their reactions

Why does Mg form Mg2+ ions and not Mg3+Why do nonmetals take on electrons

Electron Affinity A few elements GAIN electrons to form anions

Electron affinity is the energy involved when an anion loses an electron A -(g) rarr A(g) + e- EA = ΔE

EA = 141 kJΔE is ENDOthermic because O has an affinity for an e-

O [He] - ion

[He] O atom

EA = 0 kJΔE is zero for N-

due to electron-electron repulsions

Trend 5 Electron AffinityAffinity for electron increases across a period (EA becomes more negative)

Reason

Affinity decreases down a group (EA becomes less negative)

Reason

[He] N- ion

[He] N atom

Atom EAF +328 kJCl+349 kJBr+325 kJI +295 kJ

1 2 3 4 5 6 7S1

S2

S3

S4

0

50

100

150

200

250

300

350

Elec

tron

affi

nity

(kJm

ol)

Group

Period

02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

Increasing Periodic Trends

electronegativity ionization energy ionic radii electron affinityatomic radii

ionic amp atomic radii

ionization energyelectron affinity amp electronegativity

s orbitals have no planar node (l = 0) and so are spherical

p orbitals have l = 1 and have 1 planar node and so are ldquodumbbellrdquo shaped

This means d orbitals (with l = 2) have 2 planar nodes

f-orbital (quadruple dumbbell and dumbbelldouble donut)7 orbitalsWhen n = 4 When n = 4 l = 0 1 2 3 so there are 4 subshells in the shell = 0 1 2 3 so there are 4 subshells in the shell

For For l = 0 = 0 ml = 0 ---gt s subshell with single orbital = 0 ---gt s subshell with single orbital

For For l = 1 = 1 ml = -1 0 +1 ---gt p subshell with 3 orbitals = -1 0 +1 ---gt p subshell with 3 orbitals

For For l = 2 = 2 ml = -2 -1 0 +1 +2 = -2 -1 0 +1 +2 ---gt d subshell with 5 orbitals ---gt d subshell with 5 orbitals

For For l = 3 = 3 ml = -3 -2 -1 0 +1 +2 +3 = -3 -2 -1 0 +1 +2 +3 ---gt f subshell with 7 orbitals ---gt f subshell with 7 orbitals

Aufbau principleAs protons are added one by one to the nucleus to build up the elements electrons

are similarly added to these hydrogen-like orbitals

electron filling order (aufbau)

Hundrsquos Rule The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate (same energy) orbitals

Valence electrons electrons in the outermost principle quantum level of an atomThe inner electron are called core electrons

atom valence electronsCa 2N 5Br 7Arrangement of Electrons in Atoms

Shells (energy levels) = ndarr

Sublevels = ldarr

Orbitals = ml

Assigning electrons to atomsAssigning electrons to atomsElectrons generally assigned to orbitals of successively higher energyFor H atoms - Cn2 E depends only on nFor many-electron atoms energy depends on both n and l

Effective Nuclear Charge Zeff

Zeff is the nuclear charge experienced by the outermost electrons Explains why E(2s) lt E(2p)

Zeff increases across a period owing to incomplete shielding by inner electrons

Estimate Zeff by --gt [ Z - (no inner electrons) ]

Charge felt by 2s e- in Li Zeff = 3 - 2 = 1Be Zeff = 4 - 2 = 2B Zeff = 5 - 2 = 3and so on

Writing Atomic Electron Configurations

Orbital Notation Each electron is represented by a half-arrow

When electrons share an orbital opposite spins are represented by half-arrows going in opposite directions

Examples Write the orbital notation for the following atoms

N (7 e-) ____ ____ ____ ____ ____

1s 2s 2p

P ( ___ e-)

Electron Configuration Notation Write the energy level and orbital (in increasing energy) Write the number of electrons in the sublevel with a superscript

Examples Write the electron configuration notation for the following atoms

N

P

Tc

Electron Configuration Notation (Noble Gas Shorthand) Start with the noble gas ldquocloset without going overrdquo in brackets Continue normal electron configuration

Examples Write the electron configuration notation for the PREVIOUS atoms (use the noble gas shorthand)

Lewis Dot Diagrams

Valence electrons (e- in the outermost energy level) are the only electrons that are lost or gained in ordinary chemical reactions

Lewis Dot Diagrams only include electrons in valence (highest energy level) Write the letter and dots around the outside (2 on each side maximum of 8)

Examples Draw the Lewis Dot Diagram for

N Ca Cl

Ion configurations To form cations from elements remove 1 or more e- from subshell of highest n [or highest (n + l)]

P [Ne] 3s2 3p3 - 3e- ---gt P3+ [Ne] 3s2 3p0

For transition metals remove ns electrons and then (n - 1) electronsFe [Ar] 4s2 3d5

loses 2 electrons ---gt Fe2+ [Ar] 4s0 3d6

loses 3 electrons ---gt Fe3+

[Ar] 4s0 3d5

Information Contained on the Periodic Table1 Each group member has the same valence electron configuration

(these electrons primarily determine an atomrsquos chemistry)

2 The electron configuration of any representative element (see order)

3 Certain groups have special names (alkali metals halogens etc)02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

4 Metals and nonmetals are characterized by their chemical and physical properties

Periodic Trends overall change in physical property through row or group on the periodic

table These are general trends (there are exceptions)

Trend 1 Atomic Radius

Trend 2 Ionic Radius

Cations (positive ions) generally are smaller than their respective atoms

Anions (negative ions) generally are larger than their respective atoms

Trend 3 ndash Ionization Energy (IE)

Trend 4 ndash Electronegativity (en)

Redox Reactions Why do metals lose electrons in their reactions

Why does Mg form Mg2+ ions and not Mg3+Why do nonmetals take on electrons

Electron Affinity A few elements GAIN electrons to form anions

Electron affinity is the energy involved when an anion loses an electron A -(g) rarr A(g) + e- EA = ΔE

EA = 141 kJΔE is ENDOthermic because O has an affinity for an e-

O [He] - ion

[He] O atom

EA = 0 kJΔE is zero for N-

due to electron-electron repulsions

Trend 5 Electron AffinityAffinity for electron increases across a period (EA becomes more negative)

Reason

Affinity decreases down a group (EA becomes less negative)

Reason

[He] N- ion

[He] N atom

Atom EAF +328 kJCl+349 kJBr+325 kJI +295 kJ

1 2 3 4 5 6 7S1

S2

S3

S4

0

50

100

150

200

250

300

350

Elec

tron

affi

nity

(kJm

ol)

Group

Period

02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

Increasing Periodic Trends

electronegativity ionization energy ionic radii electron affinityatomic radii

ionic amp atomic radii

ionization energyelectron affinity amp electronegativity

Aufbau principleAs protons are added one by one to the nucleus to build up the elements electrons

are similarly added to these hydrogen-like orbitals

electron filling order (aufbau)

Hundrsquos Rule The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate (same energy) orbitals

Valence electrons electrons in the outermost principle quantum level of an atomThe inner electron are called core electrons

atom valence electronsCa 2N 5Br 7Arrangement of Electrons in Atoms

Shells (energy levels) = ndarr

Sublevels = ldarr

Orbitals = ml

Assigning electrons to atomsAssigning electrons to atomsElectrons generally assigned to orbitals of successively higher energyFor H atoms - Cn2 E depends only on nFor many-electron atoms energy depends on both n and l

Effective Nuclear Charge Zeff

Zeff is the nuclear charge experienced by the outermost electrons Explains why E(2s) lt E(2p)

Zeff increases across a period owing to incomplete shielding by inner electrons

Estimate Zeff by --gt [ Z - (no inner electrons) ]

Charge felt by 2s e- in Li Zeff = 3 - 2 = 1Be Zeff = 4 - 2 = 2B Zeff = 5 - 2 = 3and so on

Writing Atomic Electron Configurations

Orbital Notation Each electron is represented by a half-arrow

When electrons share an orbital opposite spins are represented by half-arrows going in opposite directions

Examples Write the orbital notation for the following atoms

N (7 e-) ____ ____ ____ ____ ____

1s 2s 2p

P ( ___ e-)

Electron Configuration Notation Write the energy level and orbital (in increasing energy) Write the number of electrons in the sublevel with a superscript

Examples Write the electron configuration notation for the following atoms

N

P

Tc

Electron Configuration Notation (Noble Gas Shorthand) Start with the noble gas ldquocloset without going overrdquo in brackets Continue normal electron configuration

Examples Write the electron configuration notation for the PREVIOUS atoms (use the noble gas shorthand)

Lewis Dot Diagrams

Valence electrons (e- in the outermost energy level) are the only electrons that are lost or gained in ordinary chemical reactions

Lewis Dot Diagrams only include electrons in valence (highest energy level) Write the letter and dots around the outside (2 on each side maximum of 8)

Examples Draw the Lewis Dot Diagram for

N Ca Cl

Ion configurations To form cations from elements remove 1 or more e- from subshell of highest n [or highest (n + l)]

P [Ne] 3s2 3p3 - 3e- ---gt P3+ [Ne] 3s2 3p0

For transition metals remove ns electrons and then (n - 1) electronsFe [Ar] 4s2 3d5

loses 2 electrons ---gt Fe2+ [Ar] 4s0 3d6

loses 3 electrons ---gt Fe3+

[Ar] 4s0 3d5

Information Contained on the Periodic Table1 Each group member has the same valence electron configuration

(these electrons primarily determine an atomrsquos chemistry)

2 The electron configuration of any representative element (see order)

3 Certain groups have special names (alkali metals halogens etc)02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

4 Metals and nonmetals are characterized by their chemical and physical properties

Periodic Trends overall change in physical property through row or group on the periodic

table These are general trends (there are exceptions)

Trend 1 Atomic Radius

Trend 2 Ionic Radius

Cations (positive ions) generally are smaller than their respective atoms

Anions (negative ions) generally are larger than their respective atoms

Trend 3 ndash Ionization Energy (IE)

Trend 4 ndash Electronegativity (en)

Redox Reactions Why do metals lose electrons in their reactions

Why does Mg form Mg2+ ions and not Mg3+Why do nonmetals take on electrons

Electron Affinity A few elements GAIN electrons to form anions

Electron affinity is the energy involved when an anion loses an electron A -(g) rarr A(g) + e- EA = ΔE

EA = 141 kJΔE is ENDOthermic because O has an affinity for an e-

O [He] - ion

[He] O atom

EA = 0 kJΔE is zero for N-

due to electron-electron repulsions

Trend 5 Electron AffinityAffinity for electron increases across a period (EA becomes more negative)

Reason

Affinity decreases down a group (EA becomes less negative)

Reason

[He] N- ion

[He] N atom

Atom EAF +328 kJCl+349 kJBr+325 kJI +295 kJ

1 2 3 4 5 6 7S1

S2

S3

S4

0

50

100

150

200

250

300

350

Elec

tron

affi

nity

(kJm

ol)

Group

Period

02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

Increasing Periodic Trends

electronegativity ionization energy ionic radii electron affinityatomic radii

ionic amp atomic radii

ionization energyelectron affinity amp electronegativity

Hundrsquos Rule The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate (same energy) orbitals

Valence electrons electrons in the outermost principle quantum level of an atomThe inner electron are called core electrons

atom valence electronsCa 2N 5Br 7Arrangement of Electrons in Atoms

Shells (energy levels) = ndarr

Sublevels = ldarr

Orbitals = ml

Assigning electrons to atomsAssigning electrons to atomsElectrons generally assigned to orbitals of successively higher energyFor H atoms - Cn2 E depends only on nFor many-electron atoms energy depends on both n and l

Effective Nuclear Charge Zeff

Zeff is the nuclear charge experienced by the outermost electrons Explains why E(2s) lt E(2p)

Zeff increases across a period owing to incomplete shielding by inner electrons

Estimate Zeff by --gt [ Z - (no inner electrons) ]

Charge felt by 2s e- in Li Zeff = 3 - 2 = 1Be Zeff = 4 - 2 = 2B Zeff = 5 - 2 = 3and so on

Writing Atomic Electron Configurations

Orbital Notation Each electron is represented by a half-arrow

When electrons share an orbital opposite spins are represented by half-arrows going in opposite directions

Examples Write the orbital notation for the following atoms

N (7 e-) ____ ____ ____ ____ ____

1s 2s 2p

P ( ___ e-)

Electron Configuration Notation Write the energy level and orbital (in increasing energy) Write the number of electrons in the sublevel with a superscript

Examples Write the electron configuration notation for the following atoms

N

P

Tc

Electron Configuration Notation (Noble Gas Shorthand) Start with the noble gas ldquocloset without going overrdquo in brackets Continue normal electron configuration

Examples Write the electron configuration notation for the PREVIOUS atoms (use the noble gas shorthand)

Lewis Dot Diagrams

Valence electrons (e- in the outermost energy level) are the only electrons that are lost or gained in ordinary chemical reactions

Lewis Dot Diagrams only include electrons in valence (highest energy level) Write the letter and dots around the outside (2 on each side maximum of 8)

Examples Draw the Lewis Dot Diagram for

N Ca Cl

Ion configurations To form cations from elements remove 1 or more e- from subshell of highest n [or highest (n + l)]

P [Ne] 3s2 3p3 - 3e- ---gt P3+ [Ne] 3s2 3p0

For transition metals remove ns electrons and then (n - 1) electronsFe [Ar] 4s2 3d5

loses 2 electrons ---gt Fe2+ [Ar] 4s0 3d6

loses 3 electrons ---gt Fe3+

[Ar] 4s0 3d5

Information Contained on the Periodic Table1 Each group member has the same valence electron configuration

(these electrons primarily determine an atomrsquos chemistry)

2 The electron configuration of any representative element (see order)

3 Certain groups have special names (alkali metals halogens etc)02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

4 Metals and nonmetals are characterized by their chemical and physical properties

Periodic Trends overall change in physical property through row or group on the periodic

table These are general trends (there are exceptions)

Trend 1 Atomic Radius

Trend 2 Ionic Radius

Cations (positive ions) generally are smaller than their respective atoms

Anions (negative ions) generally are larger than their respective atoms

Trend 3 ndash Ionization Energy (IE)

Trend 4 ndash Electronegativity (en)

Redox Reactions Why do metals lose electrons in their reactions

Why does Mg form Mg2+ ions and not Mg3+Why do nonmetals take on electrons

Electron Affinity A few elements GAIN electrons to form anions

Electron affinity is the energy involved when an anion loses an electron A -(g) rarr A(g) + e- EA = ΔE

EA = 141 kJΔE is ENDOthermic because O has an affinity for an e-

O [He] - ion

[He] O atom

EA = 0 kJΔE is zero for N-

due to electron-electron repulsions

Trend 5 Electron AffinityAffinity for electron increases across a period (EA becomes more negative)

Reason

Affinity decreases down a group (EA becomes less negative)

Reason

[He] N- ion

[He] N atom

Atom EAF +328 kJCl+349 kJBr+325 kJI +295 kJ

1 2 3 4 5 6 7S1

S2

S3

S4

0

50

100

150

200

250

300

350

Elec

tron

affi

nity

(kJm

ol)

Group

Period

02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

Increasing Periodic Trends

electronegativity ionization energy ionic radii electron affinityatomic radii

ionic amp atomic radii

ionization energyelectron affinity amp electronegativity

Orbitals = ml

Assigning electrons to atomsAssigning electrons to atomsElectrons generally assigned to orbitals of successively higher energyFor H atoms - Cn2 E depends only on nFor many-electron atoms energy depends on both n and l

Effective Nuclear Charge Zeff

Zeff is the nuclear charge experienced by the outermost electrons Explains why E(2s) lt E(2p)

Zeff increases across a period owing to incomplete shielding by inner electrons

Estimate Zeff by --gt [ Z - (no inner electrons) ]

Charge felt by 2s e- in Li Zeff = 3 - 2 = 1Be Zeff = 4 - 2 = 2B Zeff = 5 - 2 = 3and so on

Writing Atomic Electron Configurations

Orbital Notation Each electron is represented by a half-arrow

When electrons share an orbital opposite spins are represented by half-arrows going in opposite directions

Examples Write the orbital notation for the following atoms

N (7 e-) ____ ____ ____ ____ ____

1s 2s 2p

P ( ___ e-)

Electron Configuration Notation Write the energy level and orbital (in increasing energy) Write the number of electrons in the sublevel with a superscript

Examples Write the electron configuration notation for the following atoms

N

P

Tc

Electron Configuration Notation (Noble Gas Shorthand) Start with the noble gas ldquocloset without going overrdquo in brackets Continue normal electron configuration

Examples Write the electron configuration notation for the PREVIOUS atoms (use the noble gas shorthand)

Lewis Dot Diagrams

Valence electrons (e- in the outermost energy level) are the only electrons that are lost or gained in ordinary chemical reactions

Lewis Dot Diagrams only include electrons in valence (highest energy level) Write the letter and dots around the outside (2 on each side maximum of 8)

Examples Draw the Lewis Dot Diagram for

N Ca Cl

Ion configurations To form cations from elements remove 1 or more e- from subshell of highest n [or highest (n + l)]

P [Ne] 3s2 3p3 - 3e- ---gt P3+ [Ne] 3s2 3p0

For transition metals remove ns electrons and then (n - 1) electronsFe [Ar] 4s2 3d5

loses 2 electrons ---gt Fe2+ [Ar] 4s0 3d6

loses 3 electrons ---gt Fe3+

[Ar] 4s0 3d5

Information Contained on the Periodic Table1 Each group member has the same valence electron configuration

(these electrons primarily determine an atomrsquos chemistry)

2 The electron configuration of any representative element (see order)

3 Certain groups have special names (alkali metals halogens etc)02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

4 Metals and nonmetals are characterized by their chemical and physical properties

Periodic Trends overall change in physical property through row or group on the periodic

table These are general trends (there are exceptions)

Trend 1 Atomic Radius

Trend 2 Ionic Radius

Cations (positive ions) generally are smaller than their respective atoms

Anions (negative ions) generally are larger than their respective atoms

Trend 3 ndash Ionization Energy (IE)

Trend 4 ndash Electronegativity (en)

Redox Reactions Why do metals lose electrons in their reactions

Why does Mg form Mg2+ ions and not Mg3+Why do nonmetals take on electrons

Electron Affinity A few elements GAIN electrons to form anions

Electron affinity is the energy involved when an anion loses an electron A -(g) rarr A(g) + e- EA = ΔE

EA = 141 kJΔE is ENDOthermic because O has an affinity for an e-

O [He] - ion

[He] O atom

EA = 0 kJΔE is zero for N-

due to electron-electron repulsions

Trend 5 Electron AffinityAffinity for electron increases across a period (EA becomes more negative)

Reason

Affinity decreases down a group (EA becomes less negative)

Reason

[He] N- ion

[He] N atom

Atom EAF +328 kJCl+349 kJBr+325 kJI +295 kJ

1 2 3 4 5 6 7S1

S2

S3

S4

0

50

100

150

200

250

300

350

Elec

tron

affi

nity

(kJm

ol)

Group

Period

02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

Increasing Periodic Trends

electronegativity ionization energy ionic radii electron affinityatomic radii

ionic amp atomic radii

ionization energyelectron affinity amp electronegativity

Tc

Electron Configuration Notation (Noble Gas Shorthand) Start with the noble gas ldquocloset without going overrdquo in brackets Continue normal electron configuration

Examples Write the electron configuration notation for the PREVIOUS atoms (use the noble gas shorthand)

Lewis Dot Diagrams

Valence electrons (e- in the outermost energy level) are the only electrons that are lost or gained in ordinary chemical reactions

Lewis Dot Diagrams only include electrons in valence (highest energy level) Write the letter and dots around the outside (2 on each side maximum of 8)

Examples Draw the Lewis Dot Diagram for

N Ca Cl

Ion configurations To form cations from elements remove 1 or more e- from subshell of highest n [or highest (n + l)]

P [Ne] 3s2 3p3 - 3e- ---gt P3+ [Ne] 3s2 3p0

For transition metals remove ns electrons and then (n - 1) electronsFe [Ar] 4s2 3d5

loses 2 electrons ---gt Fe2+ [Ar] 4s0 3d6

loses 3 electrons ---gt Fe3+

[Ar] 4s0 3d5

Information Contained on the Periodic Table1 Each group member has the same valence electron configuration

(these electrons primarily determine an atomrsquos chemistry)

2 The electron configuration of any representative element (see order)

3 Certain groups have special names (alkali metals halogens etc)02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

4 Metals and nonmetals are characterized by their chemical and physical properties

Periodic Trends overall change in physical property through row or group on the periodic

table These are general trends (there are exceptions)

Trend 1 Atomic Radius

Trend 2 Ionic Radius

Cations (positive ions) generally are smaller than their respective atoms

Anions (negative ions) generally are larger than their respective atoms

Trend 3 ndash Ionization Energy (IE)

Trend 4 ndash Electronegativity (en)

Redox Reactions Why do metals lose electrons in their reactions

Why does Mg form Mg2+ ions and not Mg3+Why do nonmetals take on electrons

Electron Affinity A few elements GAIN electrons to form anions

Electron affinity is the energy involved when an anion loses an electron A -(g) rarr A(g) + e- EA = ΔE

EA = 141 kJΔE is ENDOthermic because O has an affinity for an e-

O [He] - ion

[He] O atom

EA = 0 kJΔE is zero for N-

due to electron-electron repulsions

Trend 5 Electron AffinityAffinity for electron increases across a period (EA becomes more negative)

Reason

Affinity decreases down a group (EA becomes less negative)

Reason

[He] N- ion

[He] N atom

Atom EAF +328 kJCl+349 kJBr+325 kJI +295 kJ

1 2 3 4 5 6 7S1

S2

S3

S4

0

50

100

150

200

250

300

350

Elec

tron

affi

nity

(kJm

ol)

Group

Period

02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

Increasing Periodic Trends

electronegativity ionization energy ionic radii electron affinityatomic radii

ionic amp atomic radii

ionization energyelectron affinity amp electronegativity

loses 3 electrons ---gt Fe3+

[Ar] 4s0 3d5

Information Contained on the Periodic Table1 Each group member has the same valence electron configuration

(these electrons primarily determine an atomrsquos chemistry)

2 The electron configuration of any representative element (see order)

3 Certain groups have special names (alkali metals halogens etc)02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

4 Metals and nonmetals are characterized by their chemical and physical properties

Periodic Trends overall change in physical property through row or group on the periodic

table These are general trends (there are exceptions)

Trend 1 Atomic Radius

Trend 2 Ionic Radius

Cations (positive ions) generally are smaller than their respective atoms

Anions (negative ions) generally are larger than their respective atoms

Trend 3 ndash Ionization Energy (IE)

Trend 4 ndash Electronegativity (en)

Redox Reactions Why do metals lose electrons in their reactions

Why does Mg form Mg2+ ions and not Mg3+Why do nonmetals take on electrons

Electron Affinity A few elements GAIN electrons to form anions

Electron affinity is the energy involved when an anion loses an electron A -(g) rarr A(g) + e- EA = ΔE

EA = 141 kJΔE is ENDOthermic because O has an affinity for an e-

O [He] - ion

[He] O atom

EA = 0 kJΔE is zero for N-

due to electron-electron repulsions

Trend 5 Electron AffinityAffinity for electron increases across a period (EA becomes more negative)

Reason

Affinity decreases down a group (EA becomes less negative)

Reason

[He] N- ion

[He] N atom

Atom EAF +328 kJCl+349 kJBr+325 kJI +295 kJ

1 2 3 4 5 6 7S1

S2

S3

S4

0

50

100

150

200

250

300

350

Elec

tron

affi

nity

(kJm

ol)

Group

Period

02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

Increasing Periodic Trends

electronegativity ionization energy ionic radii electron affinityatomic radii

ionic amp atomic radii

ionization energyelectron affinity amp electronegativity

Trend 4 ndash Electronegativity (en)

Redox Reactions Why do metals lose electrons in their reactions

Why does Mg form Mg2+ ions and not Mg3+Why do nonmetals take on electrons

Electron Affinity A few elements GAIN electrons to form anions

Electron affinity is the energy involved when an anion loses an electron A -(g) rarr A(g) + e- EA = ΔE

EA = 141 kJΔE is ENDOthermic because O has an affinity for an e-

O [He] - ion

[He] O atom

EA = 0 kJΔE is zero for N-

due to electron-electron repulsions

Trend 5 Electron AffinityAffinity for electron increases across a period (EA becomes more negative)

Reason

Affinity decreases down a group (EA becomes less negative)

Reason

[He] N- ion

[He] N atom

Atom EAF +328 kJCl+349 kJBr+325 kJI +295 kJ

1 2 3 4 5 6 7S1

S2

S3

S4

0

50

100

150

200

250

300

350

Elec

tron

affi

nity

(kJm

ol)

Group

Period

02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

Increasing Periodic Trends

electronegativity ionization energy ionic radii electron affinityatomic radii

ionic amp atomic radii

ionization energyelectron affinity amp electronegativity

EA = 0 kJΔE is zero for N-

due to electron-electron repulsions

Trend 5 Electron AffinityAffinity for electron increases across a period (EA becomes more negative)

Reason

Affinity decreases down a group (EA becomes less negative)

Reason

[He] N- ion

[He] N atom

Atom EAF +328 kJCl+349 kJBr+325 kJI +295 kJ

1 2 3 4 5 6 7S1

S2

S3

S4

0

50

100

150

200

250

300

350

Elec

tron

affi

nity

(kJm

ol)

Group

Period

02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

Increasing Periodic Trends

electronegativity ionization energy ionic radii electron affinityatomic radii

ionic amp atomic radii

ionization energyelectron affinity amp electronegativity

Atom EAF +328 kJCl+349 kJBr+325 kJI +295 kJ

1 2 3 4 5 6 7S1

S2

S3

S4

0

50

100

150

200

250

300

350

Elec

tron

affi

nity

(kJm

ol)

Group

Period

02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

Increasing Periodic Trends

electronegativity ionization energy ionic radii electron affinityatomic radii

ionic amp atomic radii

ionization energyelectron affinity amp electronegativity

02_29

1H

3Li

11Na

19K

37Rb

55Cs

87Fr

4Be

12Mg

20Ca

38Sr

56Ba

88Ra

21Sc

39Y

57La

89Acdagger

22Ti

40Zr

72Hf

104Unq

23V

41Nb

73Ta

105Unp

24Cr

42Mo

74W

106Unh

25Mn

43Tc

75Re

107Uns

26Fe

44Ru

76Os

108Uno

27Co

45Rh

77Ir

109Une

110Uun

111Uuu

28Ni

46Pd

78Pt

29Cu

47Ag

79Au

30Zn

3 4 5 6 7 8 9 10 11 12

48Cd

80Hg

31Ga

49In

81Tl

5B

13Al

32Ge

50Sn

82Pb

6C

14Si

33As

51Sb

83Bi

7N

15P

34Se

52Te

84Po

8O

16S

9F

17Cl

35Br

53I

85At

10Ne

18Ar

36Kr

54Xe

86Rn

2He

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

1A

2A

Transition metals

3A 4A 5A 6A 7A

8A1

2 13 14 15 16 17

18

Alk

ali m

etal

s

Alkalineearth metals Halogens

Noblegases

Lanthanides

dagger Actinides

Increasing Periodic Trends

electronegativity ionization energy ionic radii electron affinityatomic radii

ionic amp atomic radii

ionization energyelectron affinity amp electronegativity