Chapter 5 - Electronic Structure and Periodic Trends Electromagnetic Radiation Atomic Spectra and...

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

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.

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

Quantum

• The energy of an electron in an atom is quantized: it exists only in certain fixed quantities, rather than being continuous.

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

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

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)

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

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.

Atomic Spectra and Energy Levels

energy

Light Energy and Photons

Light is a stream of small

particles called photons that

• Energy related to their frequency. Using Plank’s constant (h)

E = h• High energy with a high frequency.

Spectrum

White light that passes

through a prism

• is separated into all colors called a continuous spectrum.

• Gives the colors of a rainbow.

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.

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

The line spectra of

several elements.

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.”

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”)

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.

Publishing as Benjamin Cummings

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.

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.

Energy Emitted

n = 5 → n = 2

n = 3 → n = 2

Bohr’s Model of the atom

• Quantized, localized electrons.

In each of the following energy level changes,indicate if energy is A) absorbed B) emitted C) not changed

Learning Check

Energy Levels, Sublevels, and Orbitals

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

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).

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.

Number of Sublevels

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

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.

Note: the designation of “n” in each type of sublevel. eg. 3p

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?

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.

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.

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.

Indicate the number and type of orbitals in eachof the following:

1. 4s sublevel

2. 3d sublevel

3. n = 3

Learning Check

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

Writing Orbital Diagrams and Electron Configurations

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

Energy Diagram for SublevelsMUST KNOW!!!

Quantum

Sublevel

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

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

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.

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

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

Sublevel Blocks

Period 1 Configurations

In Period 1, the first two electrons go into the

1s orbital.

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

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

Write the electron configuration for each of the following elements:

Learning Check

Electron Configuration and the Periodic Table

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

Sublevel Blocks

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.

Publishing as Benjamin Cummings

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

Quantum

Sublevel

4s is lower than 3d

5s is lower than 4d

• 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

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

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)

4s Block

3d Block

‡ The half-filled or fully filled 3d orbitals are more stable than a filled 4s and partially filled 3d.

4p Block

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

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

Periodic Trends of the Elements

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

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

A. X is the electron-dot symbol for

a) b) c)

B. X is the electron-dot symbol of

a) b) c)

Learning Check

Atomic SizeAtomic radius •Is the distance from the nucleus to the valence electrons.

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.

Atomic Radius Across a PeriodAtomic radius decreases

• Going from left to right across a period.

• As more protons increase nuclear attraction for valence electrons.

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.

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.

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.

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

Ionization Energy Review

Ionization energy

• Is the energy it takes to remove a valence electron.

Na(g) + Energy of Na+(g) + e-

Ionization

Ionization Energy

Metals have

• 1-3 valence electrons.

• Lower ionization energies.

Ionization Energy

Nonmetals have

• 5-7 valence electrons.

• Have higher ionization energies.

Ionization Energy

Nobel gases have

• Complete octets (He has two valence electrons.)

• Have the highest ionization energies in each period.

Chapter 5 - Electronic Structure and Periodic Trends Electromagnetic Radiation Atomic Spectra and...

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