Originally constructed to represent the patterns observed in the chemical properties of the elements. Mendeleev is given the most credit for the current.

Slides:



Advertisements
Similar presentations
Different Colored Fireworks
Advertisements

Chapter 7 Atomic Structure and Periodicity. Section 7.1 Electromagnetic Radiation Copyright © Cengage Learning. All rights reserved 2 Different Colored.
 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.
Chapter 7 Atomic Structure and Periodicity. Chapter 7 Table of Contents Copyright © Cengage Learning. All rights reserved Electromagnetic Radiation.
Chapter 7 Atomic Structure and Periodicity AP*. AP Learning Objectives  LO 1.5 The student is able to explain the distribution of electrons in an atom.
Objectives To understand how the principal energy levels fill with electrons in atoms beyond hydrogen To learn about valence electrons and core electrons.
Chapter 7 Atomic Structure and Periodicity. Chapter 7 Table of Contents Copyright © Cengage Learning. All rights reserved Electromagnetic Radiation.
UNIT 6: ATOMIC STRUCTURE & ADVANCED PERIODIC TRENDS.
Chapter 7 Atomic Structure and Periodicity. Chapter 7 Table of Contents Copyright © Cengage Learning. All rights reserved Electromagnetic Radiation.
By the end of this section, I will be able to:
Electron Configuration and Periodic Properties
Different Colored Fireworks
Different Colored Fireworks

Suggested Reading Pages Section 5-3
Section 3: Periodic Trends
Periodic properties of the elements
HC CHEMISTRY HC CHEMISTRY (B) Periodicity.
SCH3U Mr. Krstovic Agenda: 1) Atomic and Ionic Trends
Periodic Relationships Among the Elements
Write the Complete Electron Configuration for:
Periodic Trends Chemistry.
What is meant by the term “chemical bond”?
CHEMISTRY Trends and Configurations
Periodic Properties of Elements
Chemical Periodicity? What?
Different Colored Fireworks
Ions, Electron Affinity and Metallic Character
PERIODICITY.
Periodic Trends Section 6.3.
The History of the Periodic
The Periodic Table Chapter 8
AP CHEMISTRY Ms. Paskowski
Questions to Consider Why do we get colors?
Section 6.3 Practice & Assessment Problems pgs
7.1 Development of The Periodic Table
Classification of Elements and Periodic Trends
LO 1.5 The student is able to explain the distribution of electrons in an atom or ion based upon data. (Sec 7.12) LO 1.6 The student is able to analyze.
5.3 Electron Configuration & Periodic Properties
Valence Electrons Highest or outermost energy level electrons
Periodic Trends Section 6.3.
Periodic Trends.
Quantum Theory & Periodicity
Different Colored Fireworks
Different Colored Fireworks
Different Colored Fireworks
Different Colored Fireworks
Periodic Trends Section 6.3.
Topic 2 & 3 :Atomic Structure & Periodicity Modified from Scheffler
Periodic Trends: Periodic Properties
Different Colored Fireworks
Different Colored Fireworks
III. Periodic Trends (p )
TRENDS IN THE PERIODIC TABLE.
Different Colored Fireworks
Orbitals Electron Configurations Orbitals.
Different Colored Fireworks
Different Colored Fireworks
Different Colored Fireworks
Section 4.5—Periodicity.
Different Colored Fireworks
Periodicity AP Chemistry By Diane Paskowski.
The Periodic Table Chapter 5.
Periodic Trends.
5.3 Electron Configuration & Periodic Properties
Chapter 8: ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY
Periodic Trends.
Section 3: Periodic Trends
Different Colored Fireworks
Periodic Trends.
Presentation transcript:

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 existence and properties of still unknown elements. Copyright © Cengage Learning. All rights reserved

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. Oxygen: 1s22s22p4 Copyright © Cengage Learning. All rights reserved

Hund’s Rule The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate (same energy) orbitals. Copyright © Cengage Learning. All rights reserved

Orbital Diagram A notation that shows how many electrons an atom has in each of its occupied electron orbitals. Oxygen: 1s22s22p4 Oxygen: 1s 2s 2p Copyright © Cengage Learning. All rights reserved

Valence Electrons The electrons in the outermost principal quantum level of an atom. 1s22s22p6 (valence electrons = 8) The elements in the same group on the periodic table have the same valence electron configuration. Copyright © Cengage Learning. All rights reserved

The Orbitals Being Filled for Elements in Various Parts of the Periodic Table Copyright © Cengage Learning. All rights reserved

EXERCISE! Determine the expected electron configurations for each of the following. a) S 1s22s22p63s23p4 or [Ne]3s23p4 b) Ba [Xe]6s2 c) Eu [Xe]6s24f7 a) 16 electrons total; 1s22s22p63s23p4 or [Ne]3s23p4 b) 56 electrons total; [Xe]6s2 c) 63 electrons total; [Xe]6s24f7 Copyright © Cengage Learning. All rights reserved

Periodic Trends Ionization Energy Electron Affinity Atomic Radius

Ionization Energy Energy required to remove an electron from a gaseous atom or ion. X(g) → X+(g) + e– Mg → Mg+ + e– I1 = 735 kJ/mol (1st IE) Mg+ → Mg2+ + e– I2 = 1445 kJ/mol (2nd IE) Mg2+ → Mg3+ + e– I3 = 7730 kJ/mol *(3rd IE) *Core electrons are bound much more tightly than valence electrons.

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 generally more strongly bound from left to right on the periodic table.

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 from the nucleus.

The Values of First Ionization Energy for the Elements in the First Six Periods

CONCEPT CHECK! 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 The graph is not linear due to electron repulsions. Some exceptions include from Be to B and N to O.

Which atom would require more energy to remove an electron? Why? Na Cl CONCEPT CHECK! Which atom would require more energy to remove an electron? Why? Na Cl Cl would require more energy to remove an electron because the electron is more tightly bound due to the increase in effective nuclear charge.

Which atom would require more energy to remove an electron? Why? Li Cs CONCEPT CHECK! Which atom would require more energy to remove an electron? Why? Li Cs Li would require more energy to remove an electron because the outer electron is on average closer to the nucleus (so more tightly bound).

CONCEPT CHECK! Which has the larger second ionization energy? Why? Lithium or Beryllium Lithium has the larger second ionization energy because then a core electron is trying to be removed which will require a lot more energy than a valence electron.

Successive Ionization Energies (KJ per Mole) for the Elements in Period 3

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 going down a group.

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 levels.

Atomic Radii for Selected Atoms

Which should be the larger atom? Why? Na Cl CONCEPT CHECK! Which should be the larger atom? Why? Na Cl Na should be the larger atom because the electrons are not bound as tightly due to a smaller effective nuclear charge.

Which should be the larger atom? Why? Li Cs CONCEPT CHECK! Which should be the larger atom? Why? Li Cs Cs should be the larger atom because of the increase in orbital sizes in successive principal quantum levels (to accommodate more electrons).

Which is lower in energy? CONCEPT CHECK! Which is larger? The hydrogen 1s orbital The lithium 1s orbital Which is lower in energy? The hydrogen 1s orbital is larger because the electrons are not as tightly bound as the lithium 1s orbital (lithium has a higher effective nuclear charge and will thus draw in the inner electrons more closely). The lithium 1s orbital is lower in energy because the electrons are closer to the nucleus.

Atomic Radius of a Metal To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE

Atomic Radius of a Nonmetal To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE

EXERCISE! Arrange the elements oxygen, fluorine, and sulfur according to increasing: Ionization energy S, O, F Atomic size F, O, S Ionization Energy: S, O, F (IE increases as you move up a column and to the right across a period.) Atomic Size: F, O, S (Atomic radius increases as you move to the left across a period and down a column.)

The Periodic Table – Final Thoughts It is the number and type of valence electrons that primarily determine an atom’s chemistry. Electron configurations can be determined from the organization of the periodic table. Certain groups in the periodic table have special names. Copyright © Cengage Learning. All rights reserved

Special Names for Groups in the Periodic Table Copyright © Cengage Learning. All rights reserved

The Periodic Table – Final Thoughts 4. Basic division of the elements in the periodic table is into metals and nonmetals. Copyright © Cengage Learning. All rights reserved

Metals Versus Nonmetals Copyright © Cengage Learning. All rights reserved

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 Melting and boiling points smoothly decrease Copyright © Cengage Learning. All rights reserved