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Chapter 5 - Electronic Structure and Periodic Trends
•Electromagnetic Radiation•Atomic Spectra and Energy Levels•Energy Levels, Sublevels, & Orbitals•Orbital Diagrams & Electron Configurations•Electron Configurations / Periodic Table•Periodic Trends of the Elements
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Background
• In the late 1800’s, Classical (or Newtonian) Physics worked so well for the macroscopic world, most academics thought we had discovered all there was to know about the subject.– In classical physics (and everyday experience!),
things are continuous (going up ramp)– But… there were a few oddities about atoms and
light that didn’t quite work.
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Quantum Mechanics
• A new set of physical laws were needed to explain what experiments were revealing about atoms and light – These laws predicted ___________ (non-continuous)
behavior– These laws lead to strange predictions of behavior that
is impossible to picture, but explains countless experiments
– Our objective is to understand how electrons behave in the atom
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Quantum
• The energy of an electron in an atom is quantized: it exists only in certain fixed quantities, rather than being continuous.
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Adding Probability
• It is physically impossible to determine a quantum particle’s position and momentum accurately.
• Water can’t, but quantum particles can “flow uphill” briefly
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Electromagnetic Radiation
Electromagnetic radiation
• Is energy that travels as waves through space.
• Is described in terms of wavelength and frequency.
• Moves at the speed of light in a vacuum.
speed of light = 3.0 x 108 m/s
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Step 1 – Understand light
• All waves have:• Wavelength: (l) horizontal
distance between two corresponding points on a wave (units are usually m)
• Frequency: (n) the number of complete wavelengths that pass a stationary point in a second
• (units are usually Hz, s-1)
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Inverse Relationship of and
The inverse relationship of wavelength and
frequency means that
• Longer wavelengths have lower frequencies.
• Shorter wavelengths have higher frequencies.
• Different types of electromagnetic radiation have different wavelengths and frequencies.
= c = 3.0 x 108 m/s
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Electromagnetic SpectrumThe electromagnetic spectrum
• Arranges forms of energy from lower to higher.
• Arranges energy from longer to shorter wavelengths.
• Shows visible light with wavelengths from 700-400 nm.
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Atomic Spectra and Energy Levels
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energy
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Light Energy and Photons
Light is a stream of small
particles called photons that
have
• Energy related to their frequency. Using Plank’s constant (h)
E = h• High energy with a high frequency.
• Low energy with a low frequency.Copyright © 2008 by Pearson Education, Inc.Publishing as Benjamin Cummings
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Spectrum
White light that passes
through a prism
• is separated into all colors called a continuous spectrum.
• Gives the colors of a rainbow.
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Learning Check
The short wavelengths of the color blue are dispersed more by the molecules in the atmosphere than longer wavelengths of visible light, which is why we see the sky is blue.
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Atomic SpectrumAn atomic spectrum consists of lines
• Of different colors formed when light from a heated element passes through a prism.
• That correspond to photons of energy emitted by electrons going to lower energy levels.
spectrum
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The line spectra of
several elements.
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Bohr’s Model of the Atom
• His goal was to create a model that would explain the atomic spectrum of Hydrogen
– Postulates that there are only certain positions about the nucleus an electron can reside – certain “allowed orbitals.”
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Bohr’s Model of the Atom
•When electrons absorb energy (in the form of light), they move from a lower energy level to a higher one.
•When electrons move from a higher energy level to a lower energy level, they give off energy in the form of light (“emission”)
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Electron Energy Levels
Electrons are arranged in specific discrete energy levels that
• Are labeled n = 1, n = 2, n = 3, and so on. Quantized
• Increase in energy as n increases.
• Have the electrons with the lowest energy in the first energy level (n=1) closest to the nucleus. Copyright © 2008 by Pearson Education, Inc.
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Energy Level Changes
An electron
• Absorbs energy to “jump” to a higher energy level.
• Falls to a lower energy level by emitting energy.
• Emits color in the visible range.
• E = h Copyright © 2008 by Pearson Education, Inc.Publishing as Benjamin Cummings
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Changes in Energy LevelsExample:
An electron in the n = 3 energy level falls to the n = 2 energy level by emitting photons with energy equal to the energy difference of the two energy levels.
Example:
An electron in the n = 5 energy level moves to the n = 2 energy level by releasing the energy equal to the energy difference of the fifth and second energy levels.
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Energy Emitted
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n = 5 → n = 2
n = 3 → n = 2
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Bohr’s Model of the atom
• Quantized, localized electrons.
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In each of the following energy level changes,indicate if energy is A) absorbed B) emitted C) not changed
1.
2.
3.
Learning Check
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Energy Levels, Sublevels, and Orbitals
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Energy LevelsEnergy levels
• Are assigned quantum numbers n = 1, 2, 3, 4 and so on. Aka principal quantum numbers
• Increase in energy as the value of n increases.
• Have a maximum number of electrons equal to 2n2.
Energy level Maximum number of electrons
n = 1 2(1)2 = 2(1) = 2
n = 2 2(2)2 = 2(4) = 8
n = 3 2(3)2 = 2(9) = 18
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Electron levels and sublevels
n = 4n = 3n = 2
n = 1• Energy levels – Have a maximum number of
electrons equal to 2n2.
• Energy level Maximum electrons• n = 1 = 2*(1)2 = 2 • n = 2 = 2*(2)2 = 8 • n = 3 = 2*(3)2 = 18
Sublevels • The s sublevel has
the lowest energy within that sublevel.
• The s sublevel is followed by the p, d, and f sublevels in order of increasing energy.
• The number of sublevels is equal to the value of the principal quantum number (n).
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Sublevels and OrbitalsEach sublevel consists of a specific
number of orbitals. • An s sublevel contains one s orbital.• A p sublevel contains three p orbitals.• A d sublevel contains five d orbitals.• An f sublevel contains seven f orbitals.
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Number of Sublevels
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n = 4 --- 2(4)2 = 2(16) = 32 --- s, p, d, f (1+3+5+7) sublevels
n = 3 --- 2(3)2 = 2(9) = 18 --- s, p, d, (1+3+5) sublevels
n = 2 --- 2(2)2 = 2(4) = 8 --- s, p (1+3) sublevels
n = 1 --- 2(1)2 = 2(1) = 2 --- s (1) sublevel only
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Energy of Sublevels
In any energy level,
• The s sublevel has the lowest energy.
• The s sublevel is followed by the p, d, and f sublevels in order of increasing energy.
• Higher sublevels are possible, but only s, p, d, and f sublevels are needed to hold the electrons in the atoms known today.
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Note: the designation of “n” in each type of sublevel. eg. 3p
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OrbitalsAn orbital • Is a three-dimensional space
around a nucleus where an electron is found most of the time.
• Has a shape that represents electron density (not a path the electron follows).
• Can hold up to 2 electrons.• Contains two electrons that
must spin in opposite directions. Why?
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Orbitals
The shape of the orbital represents the electron density or the probability of finding an electron within that space.
Quantum Mechanics provides a means to mathematically describe the probability or shape of the orbital as well as other parameters of interest.
Quantum Chemistry or Quantum Physics are topics of advanced courses and make use of advanced calculus and differential equations to describe these probabilities.
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s Orbitals
An s orbital
• Has a spherical shape around the nucleus.
• Increases in size around the nucleus as the energy level n value increases.
• Is a single orbital found in each s sublevel.
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p OrbitalsA p orbital
• Has a two-lobed shape .
• Is one of three p orbitals that make up each p sublevel.
• Increases in size as the value of n increases.
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Indicate the number and type of orbitals in eachof the following:
1. 4s sublevel
2. 3d sublevel
3. n = 3
Learning Check
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The number of
1. electrons that can occupy _______________
A) 1 B) 2 C) 3
2. __ orbitals in the __ sublevel is
1A) 1 B) 2 C) 3
3. __ orbitals in the n = __ energy level is
A) 1 B) 3 C) 5
4. Electrons that can occupy the __ sublevel are
A) 2 B) 6 C) 14
Learning Check
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Writing Orbital Diagrams and Electron Configurations
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Energy levels fill with electrons
• In order of increasing energy.
• Beginning with quantum number n = 1.
• Beginning with s followed by p, d, and f.
Order of Filling
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Energy Diagram for SublevelsMUST KNOW!!!
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Quantum
Sublevel
Sublevel
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An orbital diagram shows• Orbitals as boxes in each sublevel.• Electrons in orbitals as vertical arrows.• Electrons in the same orbital with opposite
spins (up and down vertical arrows to indicate spin).
Orbital diagram for Li
1s 2s 2p filled half-filled empty
Orbital Diagrams
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Order of FillingElectrons in an atom
• Fill orbitals in sublevels of the same type with one electron until half full,
• Then pair up in the orbitals using opposite spins.
Atomic Element Orbital Diagram Number 1s 2s 2px2py2pz
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Writing Orbital Diagrams
The orbital diagram
for carbon consists of
• Two electrons in the 1s orbital.
• Two electrons in the 2s orbital.
• One electron each in two of the 2p orbitals.
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Write the orbital diagrams for
1. Nitrogen
1s 2s 2p 3s 3p
2. Oxygen
1s 2s 2p 3s 3p
3. Magnesium
1s 2s 2p 3s 3p
Learning Check
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An electron configuration
• Lists the sublevels filling with electrons in order of increasing energy.
• Uses superscripts to show the number of electrons in each sublevel.
• For neon is as follows:
sublevel number of electrons
1s2 2s2 2p6
Electron Configuration
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Sublevel Blocks
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Period 1 Configurations
In Period 1, the first two electrons go into the
1s orbital.
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Abbreviated ConfigurationsAn abbreviated configuration shows
• The symbol of the noble gas in brackets that represents completely filled sublevels.
• The remaining electrons in order of their sublevels.
Example: Chlorine has a configuration of
1s2 2s2 2p6 3s2 3p5
[Ne]The abbreviated configuration for chlorine is [Ne] 3s2 3p5
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Period 2 Configurations
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Period 3 Configurations
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1. The correct electron configuration for ________is A) B) C)
2. The correct electron configuration for ________is A) B) C)
3. The correct electron configuration for _________A)B)C)
Learning Check
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Write the electron configuration for each of the following elements:
1. Cl
2. S
3. K
Learning Check
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Electron Configuration and the Periodic Table
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Sublevel Blocks on the Periodic Table
The periodic table consists of sublevel blocks arranged in order of increasing energy.
• Groups 1A(1)-2A(2) = s level
• Groups 3A(13)-8A(18) = p level
• Groups 3B(3) to 2B(12) = d level
• Lanthanides/Actinides = f level
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Sublevel Blocks
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Using Sublevel BlocksTo write an electronconfiguration using sublevel
blocks • Locate the element on
the periodic table.• Write each sublevel block
in order going left to right across each period starting with H.
• Write the number of electrons in each block. Copyright © 2008 by Pearson Education, Inc.
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Using the periodic table, write the electronconfiguration for Silicon.
SolutionPeriod 11s block 1s2 Period 22s → 2p blocks 2s2 2p6 Period 3 3s → 3p blocks 3s2 3p2 (at Si)
Writing all the sublevel blocks in order gives:
1s2 2s2 2p63s2 3p2
Writing Electron Configurations
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Quantum
Sublevel
Sublevel
Note:
4s is lower than 3d
5s is lower than 4d
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• The 4s orbital has a lower energy that the 3d orbitals.
• In potassium 19K, the last electron enters the 4s orbital, not the 3d
1s 2s 2p 3s 3p 4s 3d
18Ar 1s2 2s2 2p6 3s2 3p6
19K 1s2 2s2 2p6 3s2 3p6 4s1
20Ca 1s2 2s2 2p6 3s2 3p6 4s2
21Sc 1s2 2s2 2p6 3s2 3p6 4s2 3d1
22Ti 1s2 2s2 2p6 3s2 3p6 4s2 3d2
Electron Configurations d Sublevel
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Using the periodic table, write the electronconfiguration for Manganese 7B(7).
Solution Mn is in Period 4Period 11s block 1s2 Period 22s → 2p blocks 2s2 2p6 Period 3 3s → 3p blocks 3s2 3p6
Period 44s → 3d blocks 4s2 3d5 (at Mn)
Writing all the sublevel blocks in order gives: 1s2 2s2 2p63s2 3p6 4s2 3d5
Writing Electron Configurations
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Using the periodic table, write the electronconfiguration for Iodine 7A(17).
Solution I is in Period 5Period 11s block 1s2 Period 22s → 2p blocks 2s2 2p6 Period 3 3s → 3p blocks 3s2 3p6
Period 44s → 3d → 4p blocks 4s2 3d10 4p6
Period 55s → 4d → 5p blocks 5s2 4d10 5p5
Writing all the sublevel blocks in order gives:1s2 2s2 2p63s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5
(iodine)
Writing Electron Configurations
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4s Block
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3d Block
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‡
‡
‡ The half-filled or fully filled 3d orbitals are more stable than a filled 4s and partially filled 3d.
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4p Block
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1. The last two sublevels blocks in the electron configuration for Co 8B(9) Period 4 are
A) B) C)
2. The last three sublevel blocks in the electron configuration for Sn 4A(14) Period 5 are
A) B) C)
Learning Check
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Learning CheckGive the symbol of the element that has
1. [Ar]4s2 3d6
2. Four 3p electrons
3. Two electrons in the 4d sublevel
4. The element that has the electron configuration
1s2 2s2 2p6 3s2 3p6 4s2 3d2
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Periodic Trends of the Elements
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Valence ElectronsThe valence electrons • Determine the chemical properties of the elements.• Are the electrons in the s and p sublevels in the
highest energy level.• Are related to the Group number of the element.
Example: Phosphorus has 5 valence electrons 5 valence electrons
P Group 5A(15) 1s22s22p6 3s23p3
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All the elements in a group have the same number
of valence electrons.
Example:
Elements in Group 2A(2) have two (2) valence
electrons.
Be 1s2 2s2
Mg 1s2 2s2 2p6 3s2
Ca 1s2 2s2 2p6 3s2 3p6 4s2
Sr 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
Group Number and Valence Electrons
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A. X is the electron-dot symbol for
a) b) c)
B. X is the electron-dot symbol of
a) b) c)
Learning Check
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Atomic SizeAtomic radius •Is the distance from the nucleus to the valence electrons.
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Atomic Radius Within A Group
Atomic radius increases
• Going down each group of representative elements.
• As the number of energy levels increases.
• Ionization energy decreases.
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Atomic Radius Across a PeriodAtomic radius decreases
• Going from left to right across a period.
• As more protons increase nuclear attraction for valence electrons.
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Sizes of Metal Atoms and Ions
A positive ion
• Has lost its valence electrons.
• Is smaller (about half the size) than its corresponding metal atom.
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Size of Sodium IonThe sodium ion Na+
• Forms when the Na atom loses one electron from the 3rd energy level.
• Is smaller than a Na atom.
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Sizes of Nonmetal Atoms and Ions
A negative ion
• Has a complete octet.
• Increases the number of valence electrons.
• Is larger (about twice the size) than its corresponding metal atom.
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Size of Fluoride IonThe fluoride ion F-
• Forms when a valence electron is added.• Has increased repulsions due to the added
valence electron.• Is larger than F atom
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Ionization Energy Review
Ionization energy
• Is the energy it takes to remove a valence electron.
Na(g) + Energy of Na+(g) + e-
Ionization
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+
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Ionization Energy
Metals have
• 1-3 valence electrons.
• Lower ionization energies.
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Ionization Energy
Nonmetals have
• 5-7 valence electrons.
• Have higher ionization energies.
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Ionization Energy
Nobel gases have
• Complete octets (He has two valence electrons.)
• Have the highest ionization energies in each period.
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