chapter 9 chapter electrons in atoms openingwhitener/courses/sp2011/113_sp2011_chap… · electrons...

Roy Kennedy

Massachusetts Bay Community College

Wellesley Hills, MA

Introductory Chemistry, 3rd Edition

Nivaldo Tro

Chapter 9

Electrons in Atoms

and the

Periodic Table

2009, Prentice Hall

Chapter

Opening

blimp

9.1 Blimps, Balloons, and Models of the Atom

9.2 Light, Electromagnetic Radiation

9.3 The Electromagnetic Spectrum

9.4 The Bohr Model: Atoms with Orbits

9.5 The Quantum Mechanical Model: Atoms with Orbitals

9.6 Quantum Mechanical Orbitals

9.7 Electron Configurations and the Periodic Table

9.8 The Explanatory Power of the Quantum Mechanical

Model

9.9 Periodic Trends: Atomic Size, Ionization Energy, and

Metallic Character

2Tro's "Introductory Chemistry",

Chapter 9

OUTLINE

• Hydrogen blimps may burn up

• Helium blimps can not burn

• Why is H reactive, but He inert?


Chapter 9

9.1 Blimps, Balloons, and

Models of the Atom


Chapter 9

9.2 Light, Electromagnetic

Radiation


Chapter 9


Chapter 9

Classical View of the Universe• Uniuverse is made of matter or energy.

• Matter has mass and volume.

Energy doesn’t have mass and volume.

• Matter is composed of particles.

• Energy, should not be composed of particles.

• All energy has in common that it travels in waves.


Chapter 9

The Nature of Light—Its Wave

Nature

• Light is a the forms of energy. Called

electromagnetic radiation made of waves.

• Electromagnetic radiation moves through

space like waves move across the surface of a

pond

8

Speed of Energy Transmission

Fireworks contrasting

the speed of sound = 340 m/s

With

The speed of light = 3.00 x 108 m/s


Chapter 9

Electromagnetic Waves• Wave characteristics:

Wave speed,

Height (amplitude),

Wavelength,

Number of wave peaks that pass in a given time = frequency.

• All electromagnetic waves move through space at the same, constant speed.

3.00 x 108 meters per second in a vacuum = The speed of light, c.


Chapter 9

Characterizing Waves• The amplitude is the height of the wave = intensity.

• The wavelength (l) is a measure of the distance of a repeat unit.

The distance from one crest to the next.

Usually measured in nanometers.

1 nm = 1 x 10-9 m


Chapter 9

Electromagnetic Waves

Figure 9.1 showing a wave


Chapter 9

Characterizing Waves

• The frequency (n) = number of waves that pass a point in a given period of time.

The number of waves = number of cycles.

Units are hertz (Hz), or cycles/s = s-1.

• The total energy is proportional to the amplitude and frequency of the waves.

The larger the wave amplitude, the more force it has.

The more frequently the waves strike, the more total force there is.

13

The Electromagnetic Spectrum• Light passed through a prism is separated into all its

colors called a continuous spectrum.

Figure 9.2 showing a continuous

spectrum


Chapter 9

Color

• The color of light is determined by its wavelength.

Or frequency.

• White light is a mixture of all the colors of visible light.

A spectrum.

RedOrangeYellowGreenBlueViolet.

• The observed color is predominantly the colors reflected al other

are absorbed.

9.3 The Electromagnetic

Spectrum


Chapter 9


Chapter 9

Electromagnetic Spectrum

Figure 9.4 the electromagnetic

spectrum


Chapter 9

Particles of Light

• Scientists in the early 20th century showed that

electromagnetic radiation was composed of

particles we call photons.

Max Planck and Albert Einstein..

• Each wavelength of light has photons that have

a different amount of energy.


Chapter 9

The Electromagnetic Spectrum and

Photon Energy

• Short wavelength light has photons with high energy.

• High frequency light has photons with high energy.

• High-energy electromagnetic radiation can potentially damage biological molecules.

Ionizing radiation.

9.4 The Bohr Model: Atoms with

Orbits


Chapter 9


Chapter 9

Light’s Relationship to Matter

• Atoms can acquire extra energy, but

they must eventually release it,

usually in the form of light.

• However, atoms emit only very

specific wavelengths which can be

used to identify the element.

Figure 9.6 Neon light

Figure 9.7 emission

tubes


Chapter 9

Emission Spectrum

Figure 9.8 (part 1) a line emission

spectrum.


Chapter 9

Spectra

Figure 9.8 (part 2) comparison of a

continuous spectrum and the

emission spectrum of several

elements.


Chapter 9

A figure showing the comlimentary

nature of emission (a few lines) an

an absorption spectum. Hole

appear in the absorption spectrum

At the same wavelengths lines

appear in the emission sprectum.


Chapter 9

The Bohr Model of

the Atom• Neils Bohr developed a model of the atom to explain

how the atom changes when it undergoes energy transitions.

• Bohr’s major idea was that the energy of the atom was quantized, and energy in the atom depends of the electron’s position.

Quantized means only have very specific amounts of energy.

25

The Bohr Model of the Atom:

Electron Orbits• Electrons travel in orbits around the nucleus.

More like shells than planet orbits.

• The farther the electron is from the nucleus the more energy it has.

Figure 9.9 Bohr orbits


Chapter 9


Orbits and Energy, Continued• Each orbit has a specific amount

of energy.

• The energy of each orbit is characterized by an integer, n, is called a quantum number. The larger the integer, the greater the distance from the nucleus, and the greater the energy.


Chapter 9


Energy Transitions

• When the atom gains energy, the electron leaps to a

higher orbit.

• When the electron leaps from a higher energy orbit

to one that is closer to the nucleus, energy is emitted

as a photon of light—a quantum of energy.


Chapter 9

The Bohr Model of the Atom

Figure 9.11 model of energy

transitions. Lowers to higher by

absorption of energy of a photon.

Higher to lower with emission of a

photon.


Chapter 9


Ground and Excited States

• The lowest amount of energy corresponds to being in the n = 1 orbit called the ground state.

• Electrons in higher orbital are in an excited state.

• Excited state are unstable so they release energy to return to the ground state.


Chapter 9


Hydrogen Spectrum

• Every hydrogen atom has identical orbits,

• Distances between the orbits differ so transitions do not have the same energy.

• The emission spectrum that has many lines that are unique to hydrogen.

• Visible lines are all transitions to level n=2


Chapter 9


Hydrogen Spectrum, Continued

Figure 9.12 Correlation of orbital

transition and the wavelength of

light emitted


Chapter 9


Success and Failure

• Bohr model accurately predicts the spectrum

of hydrogen.

• It fails when applied to multi-electron atoms.

• A better theory was needed.

9.5 The Quantum Mechanical

Model: Atoms with Orbitals


Chapter 9


Chapter 9

The Quantum-Mechanical Model

of the Atom

• Erwin Schrödinger applied the

mathematics of probability and the

ideas of quantizing energy to the

physics equations that describe

waves, resulting in an equation that

predicts the probability of finding

an electron with a particular

amount of energy at a particular

location in the atom.

Photo of

Schrödinger


Chapter 9

The Quantum-Mechanical Model:

Orbitals

• The result is a map of regions in the atom

that have a particular probability for finding

the electron.

• An orbital is a region where we have a very

high probability of finding the electron

when it has a particular amount of energy.

9.6 Quantum Mechanical

Orbitals


Chapter 9


Chapter 9

Orbits vs. Orbitals

Exact Pathways vs. Probability Regions

Figure 9.13 a fixed path

Figure 9.14 A probability distribution


Chapter 9

Wave–Particle Duality

• We’ve seen that light has the characteristics of waves and particles (photons).

• Electrons also have the characteristics of both particles and waves.

• It is impossible to predict the exact path of an electron.


Chapter 9


Quantum Numbers

• 3 integers, called quantum

numbers, that quantize the

energy.

• The principal quantum

number, n, specifies the main

energy level.

Figure 9.15

Atom Energy

Levels


Chapter 9


Quantum Numbers, Continued

• Each principal energy shell has one or more subshells.

The number of subshells = the principal quantum number.

• The subshell is often identified s,p,d,f, that

corresponds to a particular shape.


Chapter 9

Shells and Subshells

Figure 9.19

Number of shells and subshells

In an atom


Chapter 9

How Does the 1s Subshell Differ

from the 2s Subshell?

Figure 9.20 Sizes and shapes of the

1s and 2s orbitals


Chapter 9

Probability Maps and Orbital Shape:

s Orbitals


Chapter 9


p Orbitals

Figure 9.21 Sizes and shapes of the three

P orbitals


Chapter 9


d Orbitals

Figure 9.22 d orbitals


Chapter 9

Subshells and Orbitals

The subshells in a shell of H all have the same energy, but for multielectron atoms the subshells have different energies.

s < p < d < f.

• Each subshell contains one or more orbitals.

s =1 orbital.

p = 3 orbitals.

d = 5 orbitals.

f = 7 orbitals.


Chapter 9


Energy Transitions

• Atoms gain or lose energy as the electron

leaps between orbitals in different energy

shells and subshells.

• The ground state of the electron is the

lowest energy orbital.

• Higher energy orbitals are excited states.


Chapter 9

The Bohr Model vs.

the Quantum-Mechanical Model

• Both predict the spectrum of hydrogen

accurately.

• Only the quantum-mechanical model

predicts the spectra of multi-electron atoms.

9.7 Electron Configurations and

the Periodic Table


Chapter 9


Chapter 9

Electron Configurations

• The distribution of electrons in the various energy

shells and subshells in its ground state is the

electron configuration.

• Maximum number of electrons each subshell can

hold.

s = 2, p = 6, d = 10, f = 14.

• We place electrons in the energy shells and

subshells in order of energy, from low energy up.

Aufbau principle.

• Figure 9.26 Energy levels of a multielectron

• atom


Chapter 9


Chapter 9

Filling an Orbital with Electrons

• Each orbital may have a maximum of 2 electrons.

Pauli Exclusion principle.

• Electrons spin on an axis and two electrons the same orbital must have opposite spins.

So their magnetic fields will cancel.


Chapter 9

Orbital Diagrams

• orbital as a square and the electrons as arrows.

The direction of the arrow represents the spin of the

electron.

Figure 9.? Possible electron

occupancies of an orbital diagram

0 or 1 or 2(paired spins)


Chapter 9

Order of Subshell Filling

in Ground State Electron Configurations

Draw the triangle of orbitals as

shown

Next, draw arrows through

the diagonals, looping back

to the next diagonal

each time.

Figure 9.27 orbital filling order

triangle


Chapter 9

Filling the Orbitals in a Subshell

with Electrons

• Energy shells fill from lowest energy to highest.

• Subshells fill from lowest energy to highest.

s → p → d → f

• Orbitals in thesame subshell have the same

energy.

• When filling a subshell, place one electron in each

before completing pairs.

Hund’s rule.


Chapter 9

Electron Configuration of Atoms

in their Ground State

• The electron configuration is a listing of the subshells in order with the number of electrons written as a superscript.

Kr = 36 electrons = 1s22s22p63s23p64s23d104p6

• A short-hand way of writing an electron configuration is to use the symbol of the previous noble gas in [] to represent all the inner electrons. Noble gases (except He) end with p subshells Ne (2p), Ar (3p), Kr (4p), Xe (5p)

Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1 = [Kr]5s1


Chapter 9

Electron Configurations

Page 293 orbital diagrams for

Li, Be, B, and C


Chapter 9

1. Atomic number from the periodic table = s

the number of protons and electrons

Mg Z = 12, so Mg has 12 protons and 12 electrons.

Example—Write the Ground State

Orbital Diagram and Electron

Configuration of Magnesium.

Electron configuration of any

atom or ion• Find the atomic number from periodic table = the number of electrons

for an atom.

• Fill the orbitals in the correct order, remember the maximum electron

per subshell. Use the triangle for the order.

• s = 2 e- p = 6 e- d = 10 e- f = 14 e-

• Place electron in orbital boxes, half fill boxes before pairing.

• When electrons are used up, write the electron configuration using a

subscript for the number of electron in a subshell


Chapter 9

Example Mg

• Mg, Z = 12.

• We need 12 electrons

Need up to 3s

• Fill the correct number of electrons

Total = 2 4 10 12, stop

• Write configuration 1s22s22p63s2 = [Ne]3s2

• .

60

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

+2 = 4e−

+6 +2 = 12e−

2e−

1s 2s 2p 3s

1s 2s 2p 3s 3p

3p

Example Manganese


Chapter 9

Mn, Z = 25, we need 25 electrons

If we fill to 4s = 20 and 3d = 30

So 3d is partially filled.1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

2 e−

+2 = 4e−

+6 +2 = 12e−

+10 = 30e−

+6 +2 = 20e−

Manganese continued


Chapter 9

1s 2s 2p 3s 3p 4s

3 d

Lay out the required subshells in order

Manganese continued


Chapter 9

1s 2s 2p 3s 3p 4s

3 d

Populate the orbitals with electrons. Notice reach 3d is half filled

2 4 10 12 18 20

(5 = 25, stop)

Write the electron configuration.

Mn = 1s22s22p63s23p64s23d5 or [Ar]4s23d5

.


Chapter 9

Practice—Write the Full Ground State Orbital

Diagram and Electron Configuration of Potassium,

Continued.

K Z = 19, therefore 19 e−

1s 2s 2p 3s 3p 4s

Therefore the electron configuration is

K = 1s22s22p63s23p64s1 or K = [Ar]4s1

Ions - cations subtract and anions

add electrons


Chapter 9

Example F- F = 9 e- and -1 means an extra e- 9 e- + 1 e- = 10 e-

Write configuration F- = 1s22s22p6 = [Ne]

1s 2s 2p

Ions Ca2+


Chapter 9

Ca has 20 e- and (+2) means 2 e- lost 20 e- – 2 e- = 18 e-

Ca2+ = 1s22s22p63s23p6 or [Ar]

1s 2s 2p 3s 3p


Chapter 9

Valence Electrons

• The electrons in all the subshells with the highest principal energy shells are called the valence electrons.

• Electrons in lower energy shells are called core electrons.

• Chemists have observed that one of the most important factors in the way an atom behaves, both chemically and physically, is the number of valence electrons.


Chapter 9

Valence Electrons, Continued

Rb = 37 electrons = 1s22s22p63s23p64s23d104p65s1

• The highest principal energy shell the 5th, therefore,

Rb has 1 valence electron and 36 core electrons.

Kr = 36 electrons = 1s22s22p63s23p64s23d104p6

• The highest principal energy shell of Kr that contains

electrons is the 4th, therefore, Kr has 8 valence

electrons and 28 core electrons.

• Core electrons are always a noble gas number of

electrons

Practice—Determine the Number and Types of

Valence Electrons in an As Atom, Continued.

As Z = 33, therefore 33 e−.

Therefore, the electron configuration is 1s22s22p63s23p64s23d104p3.

The valence electrons are in level n = 4 4s24p3 2+3 = 5 valence

electrons

.


Chapter 9

Electron Configurations and

the Periodic Table

Figure 9.28 Electron configuration

and the periodic table


Chapter 9

Electron Configurations from

the Periodic Table (Main Group)

• Same period (row) have valence electrons in the same principal energy shell.

• Across a row valence electronsgoes from 1 to 8.

• Elements in the same group (column) have the same number of valence electrons.


Chapter 9

Electron Configuration and the

Periodic Table

• Elements in the same column have similar chemical and physical properties because their valence shell electron configuration is the same.

• The number of valence electrons for the main group elements is the same as the group number.


Chapter 9

Figure 9.29 illustration the s block,

P block, d block, and f blocks of the

Periodic table.


Chapter 9

Electron Configuration from

the Periodic Table

• The inner electron configuration is the same as the

noble gas of the preceding period.

• To get the outer electron configuration from the

preceding noble gas, loop through the next period,

marking the subshells as you go, until you reach the

element.

The valence energy shell = the period number.

The d block is always one energy shell below the period

number and the f is two energy shells below.

• Figure 9.30, 9.31 – using the table to write

the electron configuration for P and As


Chapter 9


Chapter 9

Practice—Use the Periodic Table to Write the Short

Electron Configuration and Orbital Diagram for

Each of the Following and Determine the Number of

Valence Electrons.

• Na (at. no. 11).

• Te (at. no. 52).


Chapter 9

Practice—Use the Periodic Table to Write the Short

Electron Configuration and Orbital Diagram for

Each of the Following and Determine the Number of

Valence Electrons, Continued.

• Na (at. no. 11). [Ne]3s1 1 valence electron

• Te (at. no. 52). [Kr]5s24d105p4 6 valence electrons

3s

5s 5p4d


Chapter 9

The Explanatory Power of

the Quantum-Mechanical Model

• The properties of the elements are largely

determined by the number of valence electrons

they contain.

• Since elements in the same column have the same

number of valence electrons, they show similar

properties.

• Since the number of valence electrons increases

across the period, the properties vary in a regular

fashion.

Tro's "Introductory Chemistry",

Chapter 9

The Noble Gas

Electron Configuration• The noble gases have 8 valence electrons.

Except for He, which has only 2 electrons.

• non-reactive because the electron

configuration of the noble gases is especially

stable.

Figure 9.32

Similarity

In eletron

Configuration

Of the

Nobel gases

Tro's "Introductory Chemistry",

Chapter 9

Everyone Wants to Be Like a Noble Gas!

The Alkali Metals

• The alkali metals have one more

electron than the previous noble

gas.

• The alkali metals tend to lose their

extra electron giving electron

configuration of a noble gas.

Cation with a 1+ charge.

Figure 9.32

Similarity

In eletron

Configuration

Of the

Alkali metals

Group 1A


Chapter 9

81

Everyone Wants to Be Like a Noble Gas!

The Halogens• Halogens all have one fewer electron than

the next noble gas.

• Reaction with Metals - Gain an electron and attain the electron configuration of the next noble gas forming an anion with charge 1−.

• Reactions with nonmetals - share electrons with the other nonmetal so that each attains the electron configuration of a noble gas.

Figure 9.32

Similarity

In electron

Configuration

Of the

Halogens

Group 7A


Chapter 9

Everyone Wants to Be

Like a Noble Gas!

• Alkali metals are the most reactive metals.

• Halogens are the most reactive group of nonmetals.

• They are only one electron away from having a very

stable electron configuration.

The same as a noble gas.


Chapter 9

Stable Electron Configuration

and Ion Charge• Metals losing valence

electrons = noble gas.

• Nonmetals gaining valence electrons = noble gas.

Atom Atom’s

electron

config

Ion Ion’s

electron

config

Na [Ne]3s1

Na+

[Ne]

Mg [Ne]3s2 Mg

2+ [Ne]

Al [Ne]3s23p

1 Al

3+ [Ne]

O [He]2s2p4 O

2- [Ne]

F [He]2s22p

5 F

- [Ne]

Periodic Trends in the

Properties of the Elements


Chapter 9

Trends in Atomic Size

• Either volume or radius of sphere

• Down a column on the periodic table, the size of the atom increases.

Valence shell farther from nucleus.

Effective nuclear charge fairly close.

• Left to right across a period, the size of the atom decreases.

Adding electrons to same valence shell.

Effective nuclear charge increases.

Valence shell held closer.

86

Trends in Atomic Size, Continued

Figure 9.37 from book shows radii of main group atoms.


Chapter 9

Graphical plot of atomic radius showing the trend

Similar to this

plot from webelements

http://www.webelements.com/_media/periodicity/tables/line/atomic_radius.gif

http://www.webelements.com/_media/periodicity/tables/line/atomic_radius.gif


Chapter 9

Example 9.6 – Choose the

Larger Atom in Each Pair

• C or O

• Li or K

• C or Al

• Se or I?

Page 302 left margin


Chapter 9

Practice—Choose the

Larger Atom in Each Pair.

1. N or F

2. C or Ge

3. N or Al

4. Al or Ge

Answers next slide


Chapter 9

1. N or F

2. C or Ge

3. N or Al

4. Al or Ge? opposing trends

1. N or F, N is further left

2. Ge is lower

3. Al is down and to the left


Larger Atom in Each Pair, Continued.


Chapter 9

Ionization Energy

• Minimum energy needed to remove an electron

from an atom.

Gas state.

Endothermic process.

Valence electron.

M(g) + 1st IE M1+(g) + 1 e-

M+1(g) + 2nd IE M2+(g) + 1 e-

First ionization energy = energy to remove electron from

neutral atom; 2nd IE = energy to remove from +1 ion; etc.

92

Plot of 1st ionization energy trend from

schoolworkhelper.net

http://schoolworkhelper.net/wp-content/uploads/2010/08/Ionization-Energy2.jpg


Chapter 9

Trends in Ionization Energy

• As atomic radius increases, the ionization energy (IE) generally decreases.

Because the electron is closer to the nucleus.

• 1st IE < 2nd IE < 3rd IE …

• As you traverse down a column, the IE gets smaller.

Valence electron farther from nucleus.

• As you traverse left to right across a period, the IE gets larger.

Effective nuclear charge increases.


Chapter 9

Trends in Ionization Energy, Continued

95

1. Al or S

2. As or Sb, Sb is further down

1. Al or S

2. As or Sb

3. N or Si, Si is further down and left

1. Al or S

2. As or Sb

3. N or Si

4. O or Cl, opposing trends

Example—Choose the Atom in Each Pair

with the Higher First Ionization Energy

1. Al or S, Al is further left


Answers next page


Chapter 9

Practice—Choose the Atom with the

Highest Ionization Energy in Each Pair

1. Mg or P

2. Cl or Br

3. Se or Sb

4. P or Se


Chapter 9

Practice—Choose the Atom with the

Highest Ionization Energy in Each Pair,

Continued

1. Mg or P P is to the right

2. Cl or Br Cl is higher

3. Se or Sb Se is up and right

4. P or Se ? Opposing trends

98

Metallic Character• How well an element’s properties match the

general properties of a metal.• Metals:Malleable and ductile as solids.Solids are shiny, lustrous, and reflect light.Solids conduct heat and electricity.Most oxides basic and ionic.Form cations in solution.Lose electrons in reactions—oxidized.

• Nonmetals:Brittle in solid state.Solid surface is dull, nonreflective.Solids are electrical and thermal insulators.Most oxides are acidic and molecular.Form anions and polyatomic anions.Gain electrons in reactions—reduced.


Chapter 9

Metallic Character, Continued

• In general, metals are found on the left of

the periodic table and nonmetals on the

right.

• As you traverse left to right across the

period, the elements become less metallic.

• As you traverse down a column, the

elements become more metallic.


Chapter 9

Trends in Metallic Character


Chapter 9

1. Sn or Te

2. P or Sb, Sb is further down

1. Sn or Te

2. P or Sb

3. Ge or In, In is further down & left

1. Sn or Te

2. P or Sb

3. Ge or In

4. S or Br? opposing trends

1. Sn or Te, Sn is further left

Example—Choose the

More Metallic Element in Each Pair



Chapter 9


More Metallic Element in Each Pair

1. Sn or Te

2. Si or Sn

3. Br or Te

4. Se or I


Chapter 9


More Metallic Element in Each Pair,

Continued

1. Sn or Te Sn is farther left

2. Si or Sn Sn is lower

3. Br or Te Te is left and lower

4. Se or I ? Opposong trends

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