Modern Atomic Theory and the Periodic Table Chapter 10

Chapter 10 - Modern Atomic Theory and the Periodic Table 10.1 A Brief History 10.5 Atomic Structures of the First 18 Elements 10.2 Electromagnetic Radiation 10.6 Electron Structures and the Periodic Table 10.3 The Bohr Atom 10.4 Energy Levels of Electrons

A Brief History

Electromagnetic Radiation Examples light from the sun x-rays microwaves radio waves television waves radiant heat All show wavelike behavior. Each travels at the same speed in a vacuum. 3.00 x 108 m/s

Characteristics of a Wave Wavelength (λ) Light has the properties of a wave. wavelength (measured from peak to peak) wavelength (measured from trough to trough) 10.1

Frequency (n) is the number of wavelengths that pass a particular point per second. 10.1

Speed (v) is how fast a wave moves through space. 10.1

Light also exhibits the properties of a particle Light also exhibits the properties of a particle. Light particles are called photons. Both the wave model and the particle model are used to explain the properties of light.

The Electromagnetic Spectrum X-rays are part of the electromagnetic spectrum visible light is part of the electromagnetic spectrum Infrared light is part of the electromagnetic spectrum 10.2

The Bohr Atom At high temperatures or voltages, elements in the gaseous state emit light of different colors. When the light is passed through a prism or diffraction grating a line spectrum results.

These colored lines indicate that light is being emitted only at certain wavelengths. Each element has its own unique set of spectral emission lines that distinguish it from other elements. Line spectrum of hydrogen. Each line corresponds to the wavelength of the energy emitted when the electron of a hydrogen atom, which has absorbed energy falls back to a lower principal energy level. 10.3

Niels Bohr Niels Bohr, a Danish physicist, in 1912-1913 carried out research on the hydrogen atom.

The Bohr Atom An electron has a discrete energy when it occupies an orbit. Electrons revolve around the nucleus in orbits that are located at fixed distances from the nucleus. 10.4

The Bohr Atom The color of the light emitted corresponds to one of the lines of the hydrogen spectrum. When an electron falls from a higher energy level to a lower energy level a quantum of energy in the form of light is emitted by the atom. 10.4

The Bohr Atom Different lines of the hydrogen spectrum correspond to different electron energy level shifts. 10.4

The Bohr Atom Light is not emitted continuously. It is emitted in discrete packets called quanta. 10.4

The Bohr Atom E1 E2 E3 An electron can have one of several possible energies depending on its orbit. 10.4

The Bohr Atom Bohr’s methods did not succeed for heavier atoms. Bohr’s calculations succeeded very well in correlating the experimentally observed spectral lines with electron energy levels for the hydrogen atom. Bohr’s methods did not succeed for heavier atoms. More theoretical work on atomic structure was needed.

For objects the size of an electron the wavelength can be detected. In 1924 Louis De Broglie suggested that all objects have wave properties. De Broglie showed that the wavelength of ordinary sized objects, such as a baseball, are too small to be observed. For objects the size of an electron the wavelength can be detected.

In 1926 Erwin Schröedinger created a mathematical model that showed electrons as waves. Schröedinger’s work led to a new branch of physics called wave or quantum mechanics. Using Schröedinger’s wave mechanics, the probability of finding an electron in a certain region around the atom can be determined. The actual location of an electron within an atom cannot be determined.

Based on wave mechanics it is clear that electrons are not revolving around the nucleus in orbits. Instead of being located in orbits, the electrons are located in orbitals. An orbital is a region around the nucleus where there is a high probability of finding an electron.

According to Bohr the energies of electrons in an atom are quantized. Energy Levels of Electrons The wave-mechanical model of the atom also predicts discrete principal energy levels within the atom According to Bohr the energies of electrons in an atom are quantized.

As n increases, the energy of the electron increases. The first four principal energy levels of the hydrogen atom. Each level is assigned a principal quantum number n. 10.7

Each principal energy level is subdivided into sublevels. 10.7, 10.8

Within sublevels the electrons are found in orbitals. An s orbital is spherical in shape. The spherical surface encloses a space where there is a 90% probability that the electron may be found. 10.10

An atomic orbital can hold a maximum of two electrons. An electron can spin in one of two possible directions represented by ↑ or ↓. The two electrons that occupy an atomic orbital must have opposite spins. This is known as the Pauli Exclusion Principal. 10.10

A p sublevel is made up of three orbitals. Each p orbital has two lobes. Each p orbital can hold a maximum of two electrons. A p sublevel can hold a maximum of 6 electrons. 10.10

The P orbitals The three p orbitals share a common center. pz The three p orbitals share a common center. The three p orbitals point in different directions. px py 10.10

A d sublevel is made up of five orbitals. The five d orbitals all point in different directions. Each d orbital can hold a maximum of two electrons. A d sublevel can hold a maximum of 10 electrons. 10.11

Number of Orbitals in a Sublevel 10.8 10.10 10.11

Distribution of Subshells by Principal Energy Level 2p 2p 2p n = 3 3s 3p 3p 3p 3d 3d 3d 3d 3d n = 4 4s 4p 4p 4p 4d 4d 4d 4d 4d 4f 4f 4f 4f 4f 4f 4f

The Hydrogen Atom The diameter of hydrogen’s nucleus is about 10-13 cm. The diameter of hydrogen’s electron cloud is about 10-8 cm. In the ground state hydrogen’s single electron lies in the 1s orbital. Hydrogen can absorb energy and the electron will move to excited states. The diameter of hydrogen’s electron cloud is about 100,000 times greater than the diameter of its nucleus. 10.12

Atomic Structure of the First 18 Elements To determine the electronic structures of atoms, the following guidelines are used.

Pauli exclusion principle No more than two electrons can occupy one orbital 10.10

2 s orbital 1 s orbital Electrons occupy the lowest energy orbitals available. They enter a higher energy orbital only after the lower orbitals are filled. For the atoms beyond hydrogen, orbital energies vary as s<p<d<f for a given value of n. 10.10

Each orbital in a sublevel is occupied by a single electron before a second electron enters. For example, all three p orbitals must contain one electron before a second electron enters a p orbital. 10.10

Nuclear makeup and electronic structure of each principal energy level of an atom. number of protons and neutrons in the nucleus number of electrons in each sublevel 10.13

Electron Configuration Arrangement of electrons within their respective sublevels. 2p6 Number of electrons in sublevel orbitals Principal energy level Type of orbital

Orbital Filling Electrons are indicated by arrows: ↑ or ↓. In the following diagrams boxes represent orbitals. Electrons are indicated by arrows: ↑ or ↓. Each arrow direction represents one of the two possible electron spin states.

H ↑ 1s1 He ↑ ↓ 1s2 Filling the 1s Sublevel 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. ↑ ↓ 1s2

Li ↑ ↓ ↑ 1s22s1 Be ↑ ↓ ↑ ↓ 1s22s2 Filling the 2s Sublevel 1s 2s 1s 2s The 1s orbital is filled. Lithium’s third electron will enter the 2s orbital. Be ↑ ↓ The 2s orbital fills upon the addition of beryllium’s third and fourth electrons. 1s 2s ↑ ↓ 1s22s2

↑ ↓ B ↑ ↓ ↑ 1s22s22p1 C ↑ ↓ ↑ ↑ 1s22s22p2 Filling the 2p Sublevel 1s Boron has the first p electron. The three 2p orbitals have the same energy. It does not matter which orbital fills first. C 1s 2s 2p ↑ ↓ The second p electron of carbon enters a different p orbital than the first p electron so as to give carbon the lowest possible energy. ↑ ↑ 1s22s22p2

1s22s22p3 N ↑ ↓ ↑ ↑ ↑ ↑ ↓ O ↑ ↓ ↑ ↑ 1s22s22p4 1s 2s 2p 1s 2s 2p The third p electron of nitrogen enters a different p orbital than its first two p electrons to give nitrogen the lowest possible energy. ↑ ↑ ↑ 1s 2s 2p ↑ ↓ O There are four electrons in the 2p sublevel of oxygen. One of the 2p orbitals is now occupied by a second electron, which has a spin opposite to that of the first electron already in the orbital. ↑ ↓ ↑ ↑ 1s22s22p4

F ↑ ↓ ↑ ↓ ↑ ↓ ↑ 1s22s22p5 Ne ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ 1s22s22p6 2p 1s 2s 2p 1s There are five electrons in the 2p sublevel of fluorine. Two of the 2p orbitals are now occupied by a second electron, which has a spin opposite to that of the first electron already in the orbital. ↑ ↓ ↑ ↓ ↑ 1s22s22p5 2p Ne 1s 2s ↑ ↓ There are 6 electrons in the 2p sublevel of neon, which fills the sublevel. ↑ ↓ ↑ ↓ ↑ ↓ 1s22s22p6

↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ Na 1s22s22p63s1 Mg ↑ ↓ ↓ 1s22s22p63s2 Filling the 3s Sublevel ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ Na 1s22s22p63s1 1s 2s 2p 3s The 2s and 2p sublevels are filled. The next electron enters the 3s sublevel of sodium. Mg 1s 2s 2p 3s ↑ ↓ The 3s orbital fills upon the addition of magnesium’s twelfth electron. ↓ 1s22s22p63s2

Electron Filling Order

Sublevel energy level order: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d You can memorize this sequence or....

Mendeleev’s arrangement is the precursor to the modern periodic table. Electron Structures and the Periodic Table In 1869 Dimitri Mendeleev of Russia and Lothar Meyer of Germany independently published periodic arrangements of the elements based on increasing atomic masses. Mendeleev’s arrangement is the precursor to the modern periodic table.

Horizontal rows are called periods. Period numbers correspond to the highest occupied energy level. 10.14

Groups are numbered with Roman numerals. Elements with similar properties are organized in groups or families. Elements in the B groups are designated transition elements. Elements in the A groups are designated representative elements. 10.14

Elements in the U.S. B-groups Main group elements Elements in the U.S. A-groups Transition elements Elements in the U.S. B-groups Metals Elements on the left of the stair-step line Nonmetals Elements on the right of the stair-step line Metalloids Elements on the stair-step line

Elemental Symbols and the Periodic Table

The chemical behavior and properties of elements in a family are associated with the electron configuration of its elements. For A family elements the valence electron configuration is the same in each column. 10.15

With the exception of helium which has a filled s orbital, the nobles gases have filled p orbitals. 10.15

To write an electron configuration using a noble-gas core: 1. Find the highest atomic-numbered noble gas (Group 8A element) less than the atomic number of the element for which the configuration is being written 2. Write the elemental symbol of the noble gas in square brackets, followed by the remaining configuration

B 1s22s22p1 [He]2s22p1 Na 1s22s22p63s1 [Ne]3s1 Cl 1s22s22p63s23p5 The electron configuration of any of the noble gas elements can be represented by the symbol of the element enclosed in square brackets. B 1s22s22p1 [He]2s22p1 Na 1s22s22p63s1 [Ne]3s1 Cl 1s22s22p63s23p5 [Ne]3s23p5

The electron configuration of argon is 1s22s22p63s23p6 The elements after argon are potassium and calcium Instead of entering a 3d orbital, the valence electrons of these elements enter the 4s orbital. K 1s22s22p63s23p64s1 [Ar]4s1 Ca 1s22s22p63s23p6 4s2 [Ar]4s2

Exceptions to the conventional filling order: 1. d4 configurations generally do not exist Chromium (Z = 24): Systematic prediction: Cr: [Ar]4s23d4 But d4 is not likely, so promote an electron from the 4s sublevel: Cr: [Ar]4s13d5

2. d9 configurations generally do not exist Copper (Z = 29): Systematic prediction: Cu: [Ar]4s23d9 But d9 is not likely, so promote an electron from the 4s sublevel: Cu: [Ar]4s13d10

d orbital numbers are 1 less than the period number d orbital filling Arrangement of electrons according to sublevel being filled. 10.16

f orbital numbers are 2 less than the period number f orbital filling Arrangement of electrons according to sublevel being filled. 10.16

Period number corresponds with the highest energy level occupied by electrons in that period. 10.17

The elements of a family have the same outermost electron configuration except that the electrons are in different energy levels. The group numbers for the representative elements are equal to the total number of outermost electrons in the atoms of the group. 10.17

Chapter 10 - Modern Atomic Theory and the Periodic Table 10.1 A Brief History 10.2 Electromagnetic Radiation 10.3 The Bohr Atom – Niels Bohr description of the atom (electron orbitals). 10.4 Energy Levels of Electrons – Electron configuration (from the periodic table), s, p, d, and f orbitals. 10.5 Atomic structures of the First 18 Elements – Valence electrons, Representatives and Transition elements, Families names. 10.6 Electron Structures and the Periodic table – Relationship between group number and valence electrons.