chapter 9 electrons in atoms and the periodic table 2006, prentice hall helium gas he has 2 protons...

Chapter 9Electrons in Atoms and

the Periodic Table

2006, Prentice Hall

Helium gas

He has 2 protons andif neutral how manyelectrons?

2

CHAPTER OUTLINE

Waves Electromagnetic Radiation Dual Nature of Light Bohr Model of Atom Quantum Mechanical Model of Atom Electron Configuration Electron Configuration & Periodic Table Abbreviated Electron Configuration Periodic Properties

Blimps• blimps float because they are

filled with a gas that is less dense than the surrounding air

• early blimps used the gas hydrogen, however hydrogen’s flammability lead to the Hindenburg disaster

• blimps now use helium gas, which is not flammable

Why is hydrogen gas diatomic but helium gas not?Because hydrogen is reactive and helium is inert.

What makes hydrogen reactive?

Recall that when elements are arranged in order of increasing atomic number (# of protons), certain setsof properties periodically recur

-All group I elements have similar reactivity -All noble gases are inert-Alkali metals- +1; metals-Alkaline earth metals- +2; metals-Halogens- -1; nonmetals-Noble gases- 0; nonmetals

We need a theory/model that would explain why these occur:

Classical View of the Universe• since the time of the ancient Greeks, the stuff of the

physical universe has been classified as either matter or energy

• we define matter as the stuff of the universe that has mass and volume – therefore energy is the stuff of the universe that does not have

mass and volume

• we know that matter is ultimately composed of particles, and the properties of particles determine the properties we observe

• energy therefore should not be composed of particles, in fact the thing that all energy has in common is that it travels in waves

Light: Electromagnetic Radiation

• light is a form of energy

• light is one type of energy called electromagnetic radiation with electric and magnetic field components and travels in waves

Electromagnetic Waves1. velocity = c = speed of light

– its constant! = 2.997925 x 108 m/s (m•sec-1) in vacuum– all types of light energy travel at the same speed

2. amplitude = A = measure of the intensity of the wave, “brightness”– height of the wave

3. wavelength = = distance between crests– generally measured in nanometers (1 nm = 10-9 m)– same distance for troughs or nodes– determines color

4. frequency = = how many peaks pass a point in a second1. generally measured in Hertz (Hz), 2. 1 Hz = 1 wave/sec = 1 sec-1

Electromagnetic Waves

Nodes

Amplitude

Wavelength (λ) is the distance between any 2 successive crests or troughs.

9

10.1

WAVES

Frequency (nu,) is the number of waves produced per unit time.

Wavelength and frequency are inversely proportional.

Speed tells how fast waves travel through space.

As wavelength of a wave increases

its frequency decreases

inversely proportional

10

ELECTROMAGNETICRADIATION

Energy travels through space as electromagnetic radiation. This radiation takes many forms, such as sunlight, microwaves, radio waves, etc.

In a vacuum, all electromagnetic waves travel at the speed of light (3.00 x 108 m/s), and differ from each other in their frequency and wavelength.

11

ELECTROMAGNETICRADIATION

The classification of electromagnetic waves according to their frequency is called electromagnetic spectrum.

These waves range from -rays (short λ, high f) to radio waves (long λ, low f).Short

wavelength High

frequency

Long wavelength

Low frequency

12

10.2

Visible light is a small part of the EM

spectrum

X-rays have longer λ but lower than

-rays

Infrared waves have longer λ but lower

than visible light

ELECTROMAGNETICRADIATION

Particles of Light• Albert Einstein and other scientists in the early

20th century showed that wave properties do not completely explain electromagnetic radiation (EM) and showed that EM was composed of particle-like properties called photons– photons are particles of light energy

• each wavelength of light has photons that have a different amount of energy– the longer the wavelength, the lower the energy of the

photons

14

DUAL NATUREOF LIGHT

Scientists also have much evidence that light beams act as a stream of tiny particles, called photons.

A photon of red light

A photon of blue light

Red light has longer wavelength and less energy than blue light

15

DUAL NATUREOF LIGHT

Scientists, therefore, use both the wave and particle models for explaining light. This is referred to as the wave-particle nature of light.

Scientists also discovered that when atoms are energized at high temperatures or by high voltage, they can radiate light. Neon lights are an example of this property of atoms.

Light’s Relationship to Matter

• Atoms can acquire extra energy, but they must eventually release it

• When atoms emit energy, it always is released in the form of light

• However, atoms don’t emit all colors, only very specific wavelengths– in fact, the spectrum of wavelengths

can be used to identify the element

He

Hg

When an atom absorbs energy it reemits it as light

White Light Source

Emits at every wavelength(all colors) also called a continuous spectrum

emission spectrum of Hydrogen

H produces its own unique and distinctive emission spectrum

Spectra

The Bohr Model of the Atom

• The Nuclear Model of the atom does not explain how the atom can gain or lose energy

• Neils Bohr developed a model of the atom to explain the how the structure of the atom changes when it undergoes energy transitions

• Bohr’s major idea was that the energy of the atom was quantized, and that the amount of energy in the atom was related to the electron’s position in the atom– quantized means that the atom could only

have very specific amounts of energy

1885 to1962Nobel Prize in

Physics in 1922

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The Bohr Model of the AtomElectron Orbits

• the Bohr Model, 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

e-

e-

e-

e-

e-

Rank the electrons fromhighest to lowest energy

The Bohr Model of the AtomOrbits and Energy

• each orbit has a specific amount of energy

• the energy of each orbit is characterized by an integer - the larger the integer, the more energy an electron in that orbit has and the farther it is from the nucleus– the integer (whole #s), n, is

called a quantum number

Bohr orbits are like steps in a ladder. It is possible to be on one

step or another, but it is impossible to be between steps.

22

BOHR MODELOF ATOM

Neils Bohr, a Danish physicist, studied the hydrogen atom extensively, and developed a model for the atom that was able to explain the line spectrum.

Bohr’s model of the atom consisted of electrons orbiting the nucleus at different distances from the nucleus, called energy levels.

In this model, the electrons could only occupy particular energy levels, and could “jump” to higher levels by absorbing energy.

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BOHR MODELOF ATOM

The lowest energy level is called ground state, and the higher energy levels are called excited states.

When electrons absorb energy through heating or electricity, they move to higher energy levels and become excited.

energy

24

BOHR MODELOF ATOM

When excited electrons return to the ground state, energy is emitted as a photon of light is released.

The color (wavelength) of the light emitted is determined by the difference in energy between the two states (excited and ground).

Lower energy transition

give off red light

Higher energy transition give off blue light

The Bohr Model of the AtomGround and Excited States

• The lowest amount of energy hydrogen’s one electron can have corresponds to being in the n = 1 orbit –this is the ground state

• when the atom gains energy, the electron leaps to a higher energy orbit –this is the excited state

• the atom is less stable in an excited state, and so it will release the extra energy to return to the ground state– either all at once or in several steps

The Bohr Model of the AtomHydrogen Spectrum

• every hydrogen atom has identical orbits, so every hydrogen atom can undergo the same energy transitions

• however, since the distances between the orbits in an atom are not all the same, no two leaps in an atom will have the same energy– the closer the orbits are in energy, the lower the

energy of the photon emitted– lower energy photon = longer wavelength

• therefore we get an emission spectrum that has a lot of lines that are unique to hydrogen

The Bohr Model of the AtomHydrogen Spectrum

Which e- has longer wavelength and lower energy (red, violet or blue-green)?

The Bohr Model of the AtomSuccess and Failure

• the mathematics of the Bohr Model very accurately predicts the spectrum of hydrogen

• however its mathematics fails when applied to multi-electron atoms– it cannot account for electron-electron

interactions

• a better theory was needed

29

QUANTUM MECHANICALMODEL OF ATOM

In 1926 Erwin Shrödinger created a mathematical model that showed electrons as both particles and waves. This model was called the quantum mechanical model.

This model predicted electrons to be located in a probability region called orbitals.

An orbital is defined as a region around the nucleus where there is a high probability of finding an electron.

1887 to 1961Nobel Prize in

Physics in 1933

Orbits vs. OrbitalsPathways vs. Probability

Orbit – acts as a particle and follows a well-defined path

Orbital – acts as a wave and particlethus could end up anywhere in a

probability map

31

QUANTUM MECHANICALMODEL OF ATOM

Based on this model, there are discrete principal energy levels within the atom.

Principal energy levels are designated by n.

The electrons in an atom can exist in any principal energy level.

As n increases, the energy of the

electron increases

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10.7, 10.8

QUANTUM MECHANICALMODEL OF ATOM

Each principal energy level is subdivided into sublevels.

The sublevels are designated by the letterss, p, d and f.

As n increases, the number of sublevels increases.

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QUANTUM MECHANICALMODEL OF ATOM

Within the sublevels, the electrons are located in orbitals. The orbitals are also designated by the letters s, p, d and f.

The number of orbitals within the sublevels vary with their type.

s sublevel = 1 orbital

p sublevel = 3 orbitals

d sublevel = 5 orbitals

f sublevel = 7 orbitals

An orbital can hold a maximum of 2 electrons

= 2 electrons

= 6 electrons

= 10 electrons

= 14 electrons

How does the 1s Subshell Differ from the 2s Subshell

Probability Maps & Orbital Shapes Orbitals

Probability Maps & Orbital Shapep Orbitals

Probability Maps & Orbital Shaped Orbitals

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ELECTRONCONFIGURATION

The distribution of electrons into the various energy shells and subshells in an atom’s ground state is called its electron configuration

The electrons occupy the orbitals from the lowest energy level to the highest level (Aufbau Principal).

The energy of the orbitals on any level are in the following order: s < p < d < f.

Each orbital on a sublevel must be occupied by a single electron before a second electron enters (Hund’s Rule).

39

ELECTRONCONFIGURATION

Electron configurations can be written as:

2 p6

Principal energy level

Type of orbital

Number of electrons in

orbitals

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ELECTRONCONFIGURATION

Another notation, called the orbital notation is shown below:

1 s

Principal energy level

Type of orbital

Electrons in orbital with

opposing spins

Filling an Orbital with Electrons

• each orbital may have a maximum of 2 electrons with opposite spins– Pauli Exclusion Principle

• electrons spin on an axis– generating their own magnetic field

• when two electrons are in the same orbital, they must have opposite spins– so there magnetic fields will cancel

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↑ 1s2

H ↑ 1s1

Hydrogen has 1 electron. It will occupy the orbital of lowest energy which is the 1s.

He

Helium has two electrons. Both helium electrons occupy the 1s orbital with opposite spins.

↓

1s

1s

ELECTRONCONFIGURATION

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B 1s22s22p1↑↓ ↑1s

2p

Boron has the first p electron. The three 2p orbitals have the same energy. It does not matter which orbital fills first.

C

The second p electron of carbon enters a different p orbital than the first p due to Hund’s Rule.

1s22s22p2↑

2s

↑↓

1s

↑↓2s

↑↓

2p

↑↓

ELECTRONCONFIGURATION

44

2p

Ne

The last p electron for neon pairs up with the last lone electron and completely fills the 2nd energy level.

1s22s22p6↑1s

↑↓2s

↑↓ ↑↓

ELECTRONCONFIGURATION

↑↓ ↓

Na ↑ 1s22s22p6

Sodium has 11 electron. The first 10 will occupy the orbitals of energy levels 1 and 2.

3s

3s1

core electrons

valence electron

45

ELECTRON CONFIGURATION

As electrons occupy the 3rd energy level and higher, some anomalies occur in the order of the energy of the orbitals.

Knowledge of these anomalies is important in order to determine the correct electron configuration for the atoms.

46

10.15

ELECTRON CONFIG. & PERIODIC TABLE

47

ELECTRON CONFIG. & PERIODIC TABLE

The horizontal rows in the periodic table are called periods. The period number corresponds to the number of energy levels that are occupied in that atom.

The vertical columns in the periodic table are called groups or families. For the main-group elements, the group number corresponds to the number of electrons in the outermost filled energy level (valence electrons).

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ELECTRON CONFIG. & PERIODIC TABLE

One energy level

3 energy levels

4 energy levels

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ELECTRON CONFIG. & PERIODIC TABLE

1 valence electron

5 valence electrons

3 valence electrons

50

10.15

The valence electrons configuration for the elements in periods 1-3 are shown below.

Note that elements in the same group have similar electron configurations.

ELECTRON CONFIG. & PERIODIC TABLE

51

10.16

ELECTRON CONFIG. & PERIODIC TABLE

Arrangement of orbitals in the periodic table

52

10.16

ELECTRON CONFIG. & PERIODIC TABLE

d orbital numbers are 1 less than the period number

53

10.16

ELECTRON CONFIG. & PERIODIC TABLE

f orbital numbers are 2 less than the period number

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ABBREVIATEDELECTRON CONFIG.

When writing electron configurations for larger atoms, an abbreviated configuration is used.

In writing this configuration, the non-valence (core) electrons are summarized by writing the symbol of the noble gas prior to the element in brackets followed by configuration of the valence electrons.

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ABBREVIATEDELECTRON CONFIG.

K Z = 19

1s22s22p63s23p6 4s1

core electrons

valence electron[Ar] 4s1

Previous noble gas

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ABBREVIATEDELECTRON CONFIG.

Br Z = 35

1s22s22p63s23p6 4s2

core electrons

valence electrons[Ar] 4s23d104p5

3d10 4p5

Electron Configuration of As from the Periodic Table

As = [Ar]4s23d104p3

As has 5 valence electrons

As

1234567

1A

2A 3A 4A 5A 6A 7A

8A

4s2

Ar3d10

4p3

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TRENDS INPERIODIC PROPERTIES

The electron configuration of atoms are an important factor in the physical and chemical properties of the elements.

Some of these properties include: atomic size, ionization energy and metallic character.

These properties are commonly known as periodic properties and increase or decrease across a period or group, and are repeated in each successive period or group.

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ATOMIC SIZE

The size of the atom is determined by its atomic radius, which is the distance of the valence electron from the nucleus.

For each group of the representative elements, the atomic size increases going down the group, because the valence electrons from each energy level are further from the nucleus.

60

ATOMIC SIZE

61

4 p+

2e-

2e-

12 p+

2e-

8e-

2e-

Be (4p+ & 4e-)

Group IIA

Mg (12p+ & 12e-)

Ca (20p+ & 20e-)

20 p+

2e-

8e-

2e-

8e-

12

Mg

4

Be

Ca 20

62

ATOMIC SIZE

The atomic radius of the representative elements are affected by the number of protons in the nucleus (nuclear charge).

For elements going across a period, the atomic size decreases because the increased nuclear charge of each atom pulls the electrons closer to the nucleus, making it smaller.

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ATOMIC SIZE

Li (3p+ & 3e-)

Period 2

Be (4p+ & 4e-) B (5p+ & 5e-)

6 p+

2e-

4e-

C (6p+ & 6e-)

8 p+

2e-

6e-

O (8p+ & 8e-)

10 p+

2e-

8e-

Ne (10p+ & 10e-)

2e-

1e-

3 p+

2e-

2e-

4 p+

2e-

3e-

5 p+

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IONIZATIONENERGY

The ionization energy is the energy required to remove a valence electron from the atom in a gaseous state.

When an electron is removed from an atom, a cation (+ ion) with a 1+ charge is formed.

Na (g) + IE Na+ + e-

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IONIZATIONENERGY

The ionization energy decreases going down a group, because less energy is required to remove an electron from the outer shell since it is further from the nucleus.

Larger atomLess IE

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IONIZATIONENERGY

Going across a period, the ionization energy increases because the increased nuclear charge of the atom holds the valence electrons more tightly and therefore it is more difficult to remove.

68

IONIZATIONENERGY

In general, the ionization energy is low for metals and high for non-metals.

Review of ionization energies of elements in periods 2-4 indicate some anomalies to the general increasing trend.

69

IONIZATIONENERGY

These anomalies are caused by more stable electron configurations of the atoms in groups 2 (complete “s” sublevel) and group 5 (half-filled “p” sublevels) that cause an increase in their ionization energy compared to the next element.

Be 1s2 2s2

B 1s2 2s2 2p1

More stableHigher IE

N 1s2 2s2 2p3

O 1s2 2s2 2p4

More stable (1/2 filled)Higher IE

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METALLICCHARACTER

Metallic character is the ability of an atom to lose electrons easily.

This character is more prevalent in the elements on the left side of the periodic table (metals), and decreases going across a period and increases for elements going down a group.

71

METALLICCHARACTER

Most metallic elements

Least metallic elements

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Example 1:Select the element in each pair with the larger atomic radius:

K Bror

Larger due to less nuclear

charge

73

Example 2:

F Cor

Highest IE due to most

nuclear charge

Indicate the element in each set that has the higher ionization energy and explain your choice:

N

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THE END