Chapter 7

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Section 7.1

Electromagnetic Radiation

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Section 7.1

Questions to Consider

 _{Why do we get colors?}

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Section 7.1

 _{One of the ways that energy travels through space.}  _{Three characteristics:}

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Section 7.1

Characteristics

 _{Wavelength (}_{) – distance between two consecutive}

peaks or troughs in a wave.

 _{Frequency ( ) – number of waves (cycles) per}

second that pass a given point in space

 _{Speed (}_c_{) – speed of light (2.9979×10}8_m/s)





= 

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Section 7.1

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Section 7.1

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Section 7.2

The Nature of Matter

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Section 7.2

 _{Energy can be gained or lost only in whole number}

multiples of .

 _{A system can transfer energy only in whole quanta (or}

“packets”).

 _{Energy seems to have particulate properties too.}

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Section 7.2

 _{Energy is quantized.}

 _{Electromagnetic radiation is a stream of “particles” called}

photons.

 _{Planck’s constant = h = 6.626 × 10}-34_J_s

photon = = _ hc

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Section 7.2

The Photoelectric effect

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Section 7.2

 _{Energy has mass}

 _{Dual nature of light:}

 _{Electromagnetic radiation (and all matter) exhibits}

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Section 7.3

The Atomic Spectrum of Hydrogen

 _{Continuous spectrum (results when white light is passed}

through a prism) – contains all the wavelengths of visible light

 _{Line spectrum – each line corresponds to a discrete}

wavelength:

hydrogen atom.

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Section 7.3

Why is it significant that the color emitted from the hydrogen emission spectrum is not white?

How does the emission spectrum support the idea of quantized energy levels?

CONCEPT CHECK!

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Section 7.4

The Bohr Model

 _{Electron in a hydrogen atom moves around the nucleus}

only in certain allowed circular orbits.

 _{Bohr’s model gave hydrogen atom energy levels}

consistent with the hydrogen emission spectrum.

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Section 7.4

The Bohr Model

Electronic Transitions in the Bohr Model for the Hydrogen Atom

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Section 7.4

The Bohr Model

Electronic Transitions in the Bohr Model for the Hydrogen Atom

b) An Orbit-Transition Diagram, Which Accounts for the

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Section 7.4

The Bohr Model

 _{For a single electron transition from one energy level to}

another:

ΔE = change in energy of the atom (energy of the emitted photon)

n_final = integer; final distance from the nucleus

n_initial = integer; initial distance from the nucleus 18

2 2

final initial

1 1

= 2.178 10 J  

   _  _

 

E

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Section 7.4

The Bohr Model

 _{The model correctly fits the quantized energy levels of}

the hydrogen atom and postulates only certain allowed circular orbits for the electron.

 _{As the electron becomes more tightly bound, its energy}

becomes more negative relative to the zero-energy

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Section 7.4

The Bohr Model

 _{Bohr’s model is incorrect. This model only works for}

hydrogen.

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Section 7.4

The Bohr Model

What color of light is emitted when an

excited electron in the hydrogen atom falls

from:

a) n = 5 to n = 2

b) n = 4 to n = 2

c) n = 3 to n = 2

Which transition results in the longest

wavelength of light?

blue, λ = 434 nm green, λ = 486 nm

orange/red, λ = 657 nm

EXERCISE!

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Section 7.5

The Quantum Mechanical Model of the Atom

 _{We do not know the detailed pathway of an electron.}  _{Heisenberg uncertainty principle:}

 _{There is a fundamental limitation to just how precisely we}

can know both the position and momentum of a particle at a given time.

Δx = uncertainty in a particle’s position

Δ(mν) = uncertainty in a particle’s momentum h = Planck’s constant



m



4

    

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Section 7.5

Physical Meaning of a Wave Function (Ψ)

 _{The square of the function indicates the probability of}

finding an electron near a particular point in space.

 _{Probability distribution – intensity of color is used to}

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Section 7.5

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Section 7.5

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Section 7.5

Relative Orbital Size

 _{Difficult to define precisely.}  _{Orbital is a wave function.}

 _{Picture an orbital as a three-dimensional electron density}

map.

 _{Hydrogen 1}_s_orbital:

 _{Radius of the sphere that encloses 90% of the total}

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Section 7.6

Quantum Numbers

 _{Principal quantum number (}_n) – size and energy of the orbital.

 _{Angular momentum quantum number} ₍_l_{) – shape of}

atomic orbitals (sometimes called a subshell).

 Magnetic quantum number (m_l) – orientation of the

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Section 7.6

Quantum Numbers

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Section 7.6

Quantum Numbers

For principal quantum level n = 3, determine the number of allowed subshells (different values of l), and give the designation of each.

# of allowed subshells = 3

l

= 0, 3

s

l

= 1, 3

p

l

= 2, 3

d

EXERCISE!

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Section 7.6

Quantum Numbers

For l = 2, determine the magnetic quantum numbers (m_l) and the number of orbitals.

magnetic quantum numbers = –2, – 1, 0, 1, 2

number of orbitals = 5

EXERCISE!

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Section 7.7

Orbital Shapes and Energies

1s Orbital

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Section 7.7

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Section 7.7

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Section 7.7

3

d

_x2

-y2

Electron Spin

 Electron spin quantum number (m_s) – can be +½ or -½.

 _{Pauli exclusion principle - in a given atom no two}

electrons can have the same set of four quantum numbers.

 _{An orbital can hold only two electrons, and they must}

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Section 7.9

Polyelectronic Atoms

 _{Atoms with more than one electron.}  _{Electron correlation problem:}

 _{Since the electron pathways are unknown, the}

electron repulsions cannot be calculated exactly.

 _{When electrons are placed in a particular quantum level,}

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Section 7.9

Penetration Effect

 _{A 2}_s_{electron penetrates to the nucleus more than one in}

the 2p orbital.

 _{This causes an electron in a 2}_s_{orbital to be attracted to}

the nucleus more strongly than an electron in a 2p orbital.

 _{Thus, the 2}_s_{orbital is lower in energy than the 2}_p_orbitals

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Section 7.9

Orbital Energies

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Section 7.9

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Section 7.9

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Section 7.9

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Section 7.10

The History of the Periodic Table

 _{Originally constructed to represent the patterns observed}

in the chemical properties of the elements.

 _{Mendeleev is given the most credit for the current}

version of the periodic table because he emphasized how useful the periodic table could be in predicting the

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Section 7.11

The Aufbau Principle and the Periodic Table

Aufbau Principle

 _{As protons are added one by one to the nucleus to build}

up the elements, electrons are similarly added to hydrogen-like orbitals.

 _{An oxygen atom has an electron arrangement of two}

electrons in the 1s subshell, two electrons in the 2s subshell, and four electrons in the 2p subshell.

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Section 7.11

Hund’s Rule

 _{The lowest energy configuration for an atom is the one}

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Section 7.11

Orbital Diagram

 _{A notation that shows how many electrons an atom has}

in each of its occupied electron orbitals.

Oxygen: 1s2₂_s2₂_p4

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Section 7.11

Valence Electrons

 _{The electrons in the outermost principal quantum level of}

an atom.

1s2₂_s2₂_p6_{(valence electrons = 8)}

 _{The elements in the same group on the periodic table}

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Section 7.11

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Section 7.11

Determine the expected electron configurations for each of the following.

a) S

1s2_2s2_2p6_3s2_3p4_{or [Ne]3s}2_3p4

b) Ba

[Xe]6s2

c) Eu

[Xe]6s2_4f7

EXERCISE!

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Section 7.12

Periodic Trends in Atomic Properties

Periodic Trends

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Section 7.12

Ionization Energy

 _{Energy required to remove an electron from a gaseous atom or ion.}  _X(_g_{) → X}+₍_g)_{+ e}–

Mg → Mg+_{+ e}– _I

1 = 735 kJ/mol(1st IE)

Mg+_{→ Mg}2+_{+ e}– _I

2 = 1445 kJ/mol (2nd IE)

Mg2+_{→ Mg}3+_{+ e}– _I

3 = 7730 kJ/mol *(3rd IE)

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Section 7.12

Ionization Energy

 _{In general, as we go across a period from left to right, the}

first ionization energy increases.

 _Why?

 _{Electrons added in the same principal quantum level}

do not completely shield the increasing nuclear charge caused by the added protons.

 _{Electrons in the same principal quantum level are}

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Section 7.12

Ionization Energy

 _{In general, as we go down a group from top to bottom,}

the first ionization energy decreases.

 _Why?

 _{The electrons being removed are, on average, farther}

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Section 7.12

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Section 7.12

Periodic Trends in Atomic Properties

Explain why the graph of ionization energy versus atomic number (across a row) is not linear.

electron repulsions

Where are the exceptions?

some include from Be to B and N to O

CONCEPT CHECK!

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Section 7.12

Which atom would require more energy to remove an electron? Why?

Na Cl

CONCEPT CHECK!

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Section 7.12

Which atom would require more energy to remove an electron? Why?

Li Cs

CONCEPT CHECK!

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Section 7.12

Which has the larger second ionization energy? Why? Lithium or Beryllium

CONCEPT CHECK!

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Section 7.12

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Section 7.12

Electron Affinity

 _{Energy change associated with the addition of an}

electron to a gaseous atom.

 _X(_g_{) + e}–_{→ X}–₍_g)

 _{In general as we go across a period from left to right, the}

electron affinities become more negative.

 _{In general electron affinity becomes more positive in}

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Section 7.12

Atomic Radius

 _{In general as we go across a period from left to right, the}

atomic radius decreases.

 _{Effective nuclear charge increases, therefore the}

valence electrons are drawn closer to the nucleus, decreasing the size of the atom.

 _{In general atomic radius increases in going down a group.}  _{Orbital sizes increase in successive principal quantum}

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Section 7.12

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Section 7.12

Which should be the larger atom? Why?

Na Cl

CONCEPT CHECK!

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Section 7.12

Which should be the larger atom? Why?

Li Cs

CONCEPT CHECK!

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Section 7.12

Periodic Trends in Atomic Properties

Which is larger?

_{The hydrogen 1s orbital} _{The lithium 1}_s_orbital

Which is lower in energy?

_{The hydrogen 1}_s_orbital _{The lithium 1}_s_orbital

CONCEPT CHECK!

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Section 7.12

Atomic Radius of a Metal

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Section 7.12

Atomic Radius of a Nonmetal

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Section 7.12

Arrange the elements oxygen, fluorine, and sulfur

according to increasing:

 _{Ionization energy}

S, O, F

 _{Atomic size}

F, O, S

EXERCISE!

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Section 7.13

The Properties of a Group: The Alkali Metals

The Periodic Table – Final Thoughts

1. It is the number and type of valence electrons that primarily determine an atom’s chemistry.

2. Electron configurations can be determined from the organization of the periodic table.

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Section 7.13

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Section 7.13

The Periodic Table – Final Thoughts

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Section 7.13

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Section 7.13

The Alkali Metals

 _{Li, Na, K, Rb, Cs, and Fr}

 _{Most chemically reactive of the metals}

 _{React with nonmetals to form ionic solids}

 _{Going down group:}

 _{Ionization energy decreases}  _{Atomic radius increases}

 _{Density increases}