Chemistry 20H
Chemistry: The Central Science
Chapter 6 - The Electronic Structure of Atoms
Chapter 7 - Periodic Properties of the Elements
Objectives:
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
Discuss terms associated with the wave nature of light. (6.1)
Given a frequency or wavelength, identify the type of electromagnetic radiation involved. (6.1)
Perform calculations based on wavelength, frequency and energy of electromagnetic radiation. (6.1, 6.2)
Discuss the photoelectric effect and the particle nature of light. (6.2)
Understand how line spectra of elements led to the Bohr Model of the atom. (6.3)
Discuss the three postulates of the Bohr model. (6.3)
Be able to discuss generally how atoms emit energy. (6.3)
Discuss the development of the wave model of electron behavior and the Uncertainty Principle. (6.4)
Discuss how quantum mechanics is used to explain the atomic orbitals occupied by electrons. (6.5)
Discuss the relationship between orbitals and quantum numbers. (6.5, 6.6)
Discuss the relationship between electron spin and the Pauli Exclusion Principle. (6.7)
Be able to produce an electron configuration, orbital diagram and condensed electron configurations for any
atom or element. (6.8)
Use Hund’s rule to explain how electrons are placed as orbitals are filled. (6.8)
Understand the relationship between electron configuration and the Periodic Table. (6.9)
Discuss the historical development of the Periodic Table. (7.1)
Discuss the following periodic relationships:
o effective nuclear charge (7.2)
o size of atoms and ions (7.3)
o ionization energy (7.4)
o electron affinity (7.5)
Discuss the properties of metals, nonmetals and metalloids. (7.6)
Discuss group trends for metals and nonmetals. (7.7, 7.8)
Vocabulary
See Summary and Key Terms, pages 250, 251, 290 and 291
Chapter 6. Electronic Structure of Atoms
Problems with the Rutherford Model
1. Classical physics says atoms should emit light and destroy themselves - they don’t
2. Atoms can be induced to emit light, but they give off a line spectrum, rather than a continuous
spectrum.
3. Every atom gives off different colours of light.
4. No explanation of why different atoms have different properties, or the same properties.
Section 6.1 - The Wave Nature of Light
What is light?



o
o
o

The Wave Model of Light
o
o
o
o
o
o
o
o
o
o
o
o
2
Table 1 -
Visible spectral colors and the wavelengths of regions of the electromagnetic spectrum.
Color
Wavelength (x 10-7 m)
Frequency (waves/s)
ultraviolet
violet
blue
green
yellow
orange
red
infrared
Table 2 -
15 to 7.5 x 1014
7.5 to 7.1 x 1014
7.1 to 6.1 x 1014
6.1 to 5.2 x 1014
5.2 to 5.1 x 1014
5.1 to 4.6 x 1014
4.6 to 4.3 x 1014
less than 4.3 x 1014
Spectrum of electromagnetic radiation
Name of Radiation
radio waves
microwaves
infrared waves
visible light
ultraviolet light
X-rays
gamma rays
•
2.00 to 4.00
4.00 to 4.20
4.20 to 4.90
4.90 to 5.80
5.80 to 5.90
5.90 to 6.50
6.50 to 7.00
longer than 7.00
Approximate range of wavelength
a few metres and up
a few millimetres to a few metres
7.50 x 10-7 m to 10-4 metre
4.00 to 7.50 x 10-7 metre
1.0 x 10-8 to 4.00 x 10-7 metre
1 x 10-11 to 5.0 x 10-8 metre
less than 5 x 10-11 metre
we can identify light by its wavelength or frequency:
▫ a wave 5.00 x 10-7 m is

▫ a wave 6.80 x 10-7 m is

▫ a wave 2.60 x 10-5 m is

▫ a wave 7.80 x 10-10 m is

Speed of Light




o
o
o
Complete questions 6.13 to 6.18, even
3
Section 6.2 - Quantized Energy and Photons
o
The wave model of light is good, but it does not explain:
o
o
o
Blackbody Radiation
o
o
Heated solids emit radiation (blackbody radiation)
o The wavelength distribution depends on the temperature (i.e., “red hot” objects are cooler than “white
hot” objects).
o Why does wavelength or frequency depend on temperature?
Max Planck suggested a way out by assuming that energy comes in packets called quanta.
Planck proposed a relationship between energy and the frequency of light quanta:
o
where:
 E = energy in Joules
 h = Planck’s constant (6.6262 x 10-34 J·s)
 ν = frequency, in Hertz (1/s, s-1)
as energy increases, so does frequency.
o
o
E  h
Complete questions 6.21 to 6.28
The Photoelectric Effect and Photons
o
o
o
Einstein used the quantum to explain the
photoelectric effect:
o Light comes in particles, called
photons.
o The energy of each photon is
determined by Planck’s equation.
o Light shining on the surface of a metal
can cause electrons to be ejected from
the metal.
The electrons will only be ejected if the photons
have sufficient energy:
o Below the threshold frequency no
electrons are ejected.
o Above the threshold frequency, the
excess energy appears as the kinetic energy of the ejected electrons.
Light has wave-like AND particle-like properties.
Complete question 6.30
4
Section 6.3 - Line Spectra and the Bohr Model
Line Spectra
•
•
A white light source produces a continuous spectrum, like a rainbow.
When elements are excited, only a line spectrum of discrete wavelengths is observed.
Hydrogen Spectrum
Has line spectra in 3 regions of the electromagnetic spectrum:



Bohr’s Model
o
o
Niels Bohr adopted Planck’s ideas about the quantum and applied them to the
electrons around a nucleus.
Bohr’s model is based on three postulates:
o
o
o



The Energy States of the Hydrogen Atom
o
Bohr showed mathematically that
 1
1 
E   ( R H ) 2  2 
n

 f ni 
o
where RH is the Rydberg constant, 2.18  10−18 J, and ni and nf are the initial and final energy levels of the
electron.
5

this explains the line spectra of hydrogen:
o Lyman series. High energy, in the UV range. Represents energy transition from higher quanta to
ground state.
o Balmer series. Intermediate energy, in the visible range. Represents energy transition from
higher quanta to quantum 2.
o Paschen series. Low energy, in the IR range. Represents energy transition from higher quanta to
quantum 3.
Limitations of the Bohr Model
o
o
It cannot explain the spectra of atoms other than hydrogen.
However, the model introduces two important ideas:
o The energy of an electron is quantized: electrons exist only in certain energy levels described by quantum
numbers.
o Energy gain or loss is involved in moving an electron from one energy level to another.
Section 6.4 - The Wave Behavior of Matter
o
o
o
Knowing that light has a particle nature, it seems reasonable to ask whether matter has a wave nature.
This question was answered by Louis deBroglie.
Using Einstein’s and Planck’s equations, deBroglie derived:
o λ = h/mv
o The momentum, mν (mass x volume) is a particle property, whereas λ is a wave property.
What does this mean?


o
The Uncertainty Principle
o
o
Section 6.5 - Quantum Mechanics and Atomic Orbitals
o
o
Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter
could be incorporated.
It is known as quantum mechanics.
o The wave function describes the electron’s matter wave; it gives the probability of finding the electron.
o Electron density is another way of expressing probability.
o A region of high electron density is one where there is a high probability of finding an electron.
Orbitals and Quantum Numbers
o
o
Solving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies.
Each orbital describes a spatial distribution of electron density.
6
o
An orbital is described by a set of three quantum numbers:
o
Principal quantum number, n.


o
Azimuthal quantum number, l.



Value of l
Type of Orbital

o
Magnetic quantum number, ml.








In Summary
7
Section 6.6 - Representations of Orbitals
The s Orbitals
o
o
o
The p Orbitals
o
o
o
The d Orbitals
o
o
o
8
The f Orbitals
o
o
o
Section 6.7 - Many-Electron Atoms
Orbitals and Their Energies
o
Orbitals of the same energy are said to be degenerate.
Electron Spin and the Pauli Exclusion Principle
o
o
o
o
o
o
o
Section 6.8 - Electron Configurations and Orbital Diagrams
o
o
o
o
o
o
o
o
9
For any atom or ion:



For Instance:
 Iron (Fe) –
Electron configurations of period two elements
Element
Electron configuration
lithium
1s22s1
beryllium
1s22s2
boron
1s22s22p1
carbon
1s22s22p2
nitrogen
1s22s22p3
oxygen
1s22s22p4
fluorine
1s22s22p5
neon
1s22s22p6
Examples:
•
On a separate page, write the correct electron configuration for the following:
▫ S, P, As, Fe, Br, At, U, Na1+, O2-, Ne, Kr, Si, Al, Ca
Electron Promotion (Hybridization)





Element
Unhybridized
beryllium
boron
carbon
10
Hybridized
Orbital Diagrams are another way to illustrate the position of electrons. They take more room to draw, but can
give information concerning bonding. They are best learned by comparision with electron configuration:
eg. Na (11 protons, 11 electrons)



Hund’s Rule
Electron arrangements of period two elements
Element
lithium
beryllium
boron
carbon
nitrogen
oxygen
fluorine
neon
Electron configuration
1s22s1
1s22s12p1
1s22s12p2
1s22s12p3
1s22s22p3
1s22s22p4
1s22s22p5
1s22s22p6
Orbital Diagram
1s
2s
↑↓
↑↓
↑↓
↑↓
↑↓
↑↓
↑↓
↑↓
↑
↑
↑
↑
↑↓
↑↓
↑↓
↑↓
2p
↑
↑
↑
↑
↑↓
↑↓
↑↓
↑
↑
↑
↑
↑↓
↑↓
↑
↑
↑
↑
↑↓
Note the electron promotion for the elements from groups 2, 13 & 14.
Repeat the last assignment, giving the orbital diagrams for the elements.
11
Condensed Electron Configurations
o
o
o
o
o
o
o
o
Section 6.9 - Electron Configurations and the Periodic Table
o
o
o
o
o
o
o
Anomalous Electron Configurations
o
There are many elements that appear to violate the electron configuration guidelines.
o Examples:
 Chromium is [Ar]3d54s1 instead of [Ar]3d44s2.
 Copper is [Ar]3d104s1 instead of [Ar]3d94s2.
o
o
12
Chapter 7. Periodic Properties of the Elements
Section 7.1 - Development of the Periodic Table
o
The periodic table is the most significant tool that chemists use for organizing and recalling chemical facts.
o
In the first attempt Mendeleev and Meyer arranged the elements in order of increasing atomic weight.
o Certain elements were missing from this scheme. Mendeleev predicted a number of properties for
these missing elements.
o In 1886 Ge was discovered; the properties of Ge match Mendeleev’s predictions well.
o
o
In the modern periodic table, elements are arranged in order of increasing atomic number.
Elements in the same column (family) contain the same number of outer-shell electrons or valence electrons.
o Have similar chemical and physical properties.
Elements in the same row (period) have a predictable change in electrons and properties.
o
7.2 Effective Nuclear Charge (Zeff)
o
o
is the charge experienced by an electron on a many-electron atom.
o is not the same as the charge on the nucleus because of the effect of the inner electrons.
o The electron is attracted to the nucleus, but repelled by electrons that shield or screen it from the full
nuclear charge.
o This shielding is called the screening effect.
The nuclear charge experienced by an electron depends on its distance from the nucleus and the number of
electrons in the spherical volume out to the electron in question.
7.3 Sizes of Atoms and Ions
o
o
o
The apparent radius of an atom is determined by the closest distances separating the nuclei during collisions.
o This radius is the nonbonding radius.
The distance between the two nuclei in a diatomic molecule is called the bonding atomic radius.
o It is shorter than the nonbonding radius.
If the two atoms that make up the molecule are the same, then half the bond distance is called the
covalent radius of the atom.
Periodic Trends in Atomic Radii
o
o
o
o
Atomic size varies consistently through the periodic table.
o As we move down a group the atoms become larger.
o As we move across a period atoms become smaller.
There are two factors at work:
o the principal quantum number, n
o the effective nuclear charge, Zeff.
As the principal quantum number increases (i.e., we move down a group), the distance of the outermost
electron from the nucleus becomes larger. Hence the atomic radius increases.
As we move across the periodic table, the number of core electrons remains constant, however, the nuclear
charge increases. Therefore, there is an increased attraction between the nucleus and the outermost electrons.
This attraction causes the atomic radius to decrease.
13
Periodic Trends in Ionic Radii
o
o
o
o
o
o
o
Just as atomic size is periodic, ionic size is also periodic.
In general:
o Cations are smaller than their parent atoms.
o Electrons have been removed from the most spatially extended orbital.
o The effective nuclear charge has increased.
Therefore, the cation is smaller than the parent atom.
Anions are larger than their parent atoms.
o Electrons have been added to the most spatially extended orbital.
o This means total electron-electron repulsion has increased.
o Therefore, anions are larger than their parent atoms.
For ions with the same charge, ionic size increases down a group.
All the members of an isoelectronic series have the same number of electrons.
As nuclear charge increases in an isoelectronic series the ions become smaller:
 O2– > F1- > Na1+ > Mg2+ > Al3+
7.4 Ionization Energy
o
o
o
o
o
o
The ionization energy of an atom or ion is the minimum energy required to remove an electron from the ground
state of the isolated gaseous atom or ion.
The first ionization energy, I1, is the amount of energy required to remove an electron from a gaseous atom:
 Na(g)  Na1+(g) + e–
The second ionization energy, I2, is the energy required to remove the second electron from a gaseous ion:
 Na1+(g)  Na2+(g) + e–
Ionization energies for an element increase in magnitude as successive electrons are removed.
As each successive electron is removed, more energy is required to pull an electron away from an increasingly
more positive ion.
A sharp increase in ionization energy occurs when an inner-shell electron is removed. (this would occur when
an electron is removed from Na1+ or from Ca2+)
Periodic Trends in First Ionization Energies
o
Ionization energy generally increases across a period.
o As we move across a period, Zeff increases, making it more difficult to remove an electron.
o Two exceptions:
 removing the first p electron
 removing the fourth p electron.
o
Ionization energy decreases down a group.
o This means that the outermost electron is more readily removed as we go down a group.
o As the atom gets bigger, it becomes easier to remove an electron from the most spatially extended
orbital.
o Example: for the noble gases, the ionization energies follow the order He > Ne > Ar > Kr > Xe
Electron Configurations of Ions
o
These are derived from the electron configurations of elements with the required number of electrons added or
removed from the most accessible orbital.
o Li: [He]2s1
becomes
Li1+:
[He]
2
5
o F: [He]2s 2p
becomes
F1-:
[He]2s22p6 = [Ar]
o
Transition metals tend to lose the valence shell electrons first and then as many d electrons as are required to
reach the desired charge on the ion.
o Thus electrons are removed from 4s before the 3d, etc.
14
Section 7.5 - Electron Affinities
o
o
Electron affinity is the energy change when a gaseous atom gains an electron to form a gaseous ion.
Electron affinity and ionization energy measure the energy changes of opposite processes.
 Electron affinity: Cl(g) + e–  Cl1-(g)
Δ E = –349 kJ/mol
 Ionization energy: Cl(g)  Cl1+(g) + e–
Δ E = +1251 kJ/mol
o
Electron affinities do not change greatly as we move down in a group.
Section 7.6 - Metals, Nonmetals and Metalloids
o
Metallic character refers to the extent to which the element exhibits the physical and chemical properties of
metals.
o Metallic character increases down a group.
o Metallic character decreases from left to right across a period.
Metals
o
o
o
o
o
o
Metals are shiny and lustrous, as well as malleable and ductile.
Metals are solids at room temperature and have very high melting temperatures (exceptions: mercury is liquid
at room temperature; gallium and cesium melt just above room temperature).
Metals tend to have low ionization energies and tend to form cations easily.
Metals tend to be oxidized when they react.
Compounds of metals with nonmetals tend to be ionic substances.
o Metal oxides form basic ionic solids.
o Most metal oxides are basic:
 Metal oxide + water  metal hydroxide
 Na2O(s) + H2O(l)  2 NaOH(aq)
Metal oxides are able to react with acids to form salts and water:
 Metal oxide + acid  salt + water
 NiO(s) + 2 HCl(aq)  NiCl2 (aq) + H2O(l)
Nonmetals
o
o
o
Nonmetals are more diverse in their behavior than metals.
In general, nonmetals are nonlustrous, are poor conductors of heat and electricity, and exhibit lower melting
points than metals.
Seven nonmetallic elements exist as diatomic molecules under ordinary conditions:
 H2(g), N2(g), O2(g), F2(g), Cl2(g), Br2(l), I2(s)
o
When nonmetals react with metals, nonmetals tend to gain electrons:
 Metal + nonmetal  salt
 2 Al(s) + 3 Br2(l)  2 AlBr3 (s)
o
o
Compounds composed entirely of nonmetals are molecular substances.
Most nonmetal oxides are acidic:
 Nonmetal oxide + water  acid
 CO2 (g) + H2O(l)  H2CO3 (aq)
 P4O10 (s) + 6 H2O(l)  4 H3PO4 (aq)
Nonmetal oxides react with bases to form salts and water:
 Nonmetal oxide + base  salt + water
 CO2 (g) + 2 NaOH(aq)  Na2CO3 (aq) + H2O(l)
o
15
Metalloids
o
o
Metalloids have properties that are intermediate between those of metals and nonmetals.
o Example: Si has a metallic luster but it is brittle.
Metalloids have found fame in the semiconductor industry.
Section 7.7 - Group Trends for the Active Metals
o
The alkali metals (group 1A) and the alkaline earth metals (group 2A) are often called the active metals.
Group 1A: The Alkali Metals
o
o
o
o
o
o
o
o
o
The alkali metals are in Group 1A.
Alkali metals are all soft.
Their chemistry is dominated by the loss of their single s electron:
o M  M1+ + e–
Reactivity increases as we move down the group.
Alkali metals react with hydrogen to form hydrides.
o In hydrides, the hydrogen is present as H1-, called the hydride ion.
 2 M(s) + H2 (g)  2 MH(s)
Alkali metals react with water to form MOH and hydrogen gas:
 2 M(s) + 2 H2O(l)  2 MOH(aq) + H2 (g)
Alkali metals produce different oxides when reacting with O 2:
 4 Li(s) + O2 (g)  2 Li2O(s)
(oxide)
 2 Na(s) + O2 (g)  Na2O2 (s)
(peroxide)
 K(s) + O2 (g)  KO2 (s)
(superoxide)
Alkali metals emit characteristic colors when placed in a high-temperature flame.
The s electron is excited by the flame and emits energy when it returns to the ground state.
o The Na line occurs at 589 nm (yellow).
o The Li line is crimson red.
o The K line is lilac.
Group 2A: The Alkaline Earth Metals
o
o
o
Alkaline earth metals are harder and more dense than the alkali metals.
Their chemistry is dominated by the loss of two s electrons:
o M  M2+ + 2 e–
Mg(s) + Cl2(g)  MgCl2(s)
o
o 2 Mg(s) + O2(g)  2 MgO(s)
Reactivity increases down the group.
o Be does not react with water.
o Mg will only react with steam.
o Ca and the elements below it react with water at room temperature as follows:
 Ca(s) + 2 H2O(l)  Ca(OH)2 (aq) + H2 (g)
16
Section 7.8 - Group Trends for Selected Nonmetals
Hydrogen
o
o
o
o
o
o
Hydrogen is a unique element.
It most often occurs as a colorless diatomic gas, H2.
Reactions between hydrogen and nonmetals can be very exothermic:
 2 H2 (g) + 2 O2 (g)  2 H2O(l)
ΔHo = –571.7 kJ
It can either gain another electron to form the hydride ion, H 1-, or lose its electron to become H1+:
 2 Na(s) + H2 (g)  2 NaH(s)
 2 H2 (g) + O2 (g)  2 H2O(l)
H1+ is a proton.
The aqueous chemistry of hydrogen is dominated by H1+(aq).
Group 6A: The Oxygen Group
o
o
o
As we move down the group the metallic character increases.
o O2 is a gas, Te is a metalloid, Po is a metal.
There are two important forms of oxygen: O 2 and ozone, O3.
o O2 and O3 are allotropes.
o Allotropes are different forms of the same element in the same state (in this case, gaseous).
o Ozone can be prepared from oxygen:
 3 O2 (g)  2 O3 (g)
ΔHo = +284.6 kJ
o Ozone is pungent and toxic.
o Oxygen (or dioxygen, O2) is a potent oxidizing agent since the O 2– ion has a noble gas configuration.
o There are two oxidation states for oxygen: –2 (e.g., H2O) and –1 (e.g., H2O2).
Sulfur is another important member of this group.
o The most common form of sulfur is yellow S8.
o Sulfur tends to form S2– in compounds (sulfides).
Group 7A: The Halogens
o
o
o
o
o
o
o
o
Group 7A elements are known as the halogens ("salt formers").
The chemistry of the halogens is dominated by gaining an electron to form an anion:
 X2 + 2 e–  2 X1Fluorine is one of the most reactive substances known:

2F2 (g) + 2 H2O(l)  4 HF(aq) + O2 (g)
ΔH = –758.9 kJ
All halogens consist of diatomic molecules, X2.
Chlorine is the most industrially useful halogen.
The reaction between chorine and water produces hypochlorous acid (HOCl), which is used to disinfect
swimming pool water:
 Cl2 (g) + H2O(l)  HCl(aq) + HOCl(aq)
Halogens react with hydrogen to form gaseous hydrogen halide compounds:
 H2 (g) + X2  2 HX(g)
Hydrogen compounds of the halogens are all strong acids with the exception of HF.
Group 8A: The Noble Gases
o
o
o
o
The group 8A elements are known as the noble gases.
These are all nonmetals and monoatomic.
They are notoriously unreactive because they have completely filled s and p subshells.
In 1962 the first compounds of the noble gases were prepared: XeF2, XeF4, and XeF6.
17
Assignment:
Chapter 7, pages 292 to 295
Questions 14, 22, 24, 32, 34, 38, 42, 44, 54, 64, 68, 78
Hand in on Friday, June 3
18