The Periodic Table and Periodic Law

1. The Periodic Table and Periodic Law Chapter 6

2. Section 6.1 - Development of the Modern Periodic Table 1. History of the Periodic Table’s Development • In the 1700s, Lavoisier compiled a list of all the known elements of the time. • List contained 23 elements • The 1800s brought large amounts of information, including an explosion in the number of known elements. • Approximately 70 known elements • Scientists needed a way to organize knowledge about elements. • A significant step came when chemists agreed upon a method for accurately determining the atomic masses of the elements.

3. History of the Periodic Table’s Development (cont.) • John Newlands • In 1864, English chemist John Newlands (1837-1898) proposed an arrangement where elements were ordered by increasing atomic mass. • Newlands noticed when the elements were arranged by increasing atomic mass, their properties repeated every eighth element. • The pattern was called periodic because it repeats in a specific manner. • He called it the law of octaves • Unfortunately the law didn’t work for all the known elements

4. History of the Periodic Table’s Development (cont.) • Meyer, Mendeleev, & Moseley • In 1869, German chemist Lothar Meyer and Russian chemist Dmitri Mendeleev both demonstrated a connection between atomic mass and elemental properties. • Mendeleev arranged the table by increasing atomic mass, which resulted in a periodic pattern in properties, the first periodic table. • He predicted the existence and properties of undiscovered elements. • He left blank spaces where he thought undiscovered elements should go • He predicted the properties of scandium, gallium, and germanium

5. History of the Periodic Table’s Development (cont.) • Mendeleev’s table was not completely correct, some elements were not in the correct order • In 1913, English chemist Henry Moseley discovered that atoms of each element contain unique number of protons in their nuclei. • rearranged the table by increasing atomic number, and resulted in a clear periodic pattern. • Periodic repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number is called periodic law.

6. 2. The Modern Periodic Table • The modern periodic table contains boxes which contain the element's name, symbol, atomic number, and atomic mass. • Columns of elements are called groups. • Rows of elements are called periods. • Elements in groups 1,2, and 13-18 possess a wide variety of chemical and physical properties and are called the representative elements. • Elements in groups 3-12 are known as the transition metals.

7. The Modern Periodic Table (cont.) • Elements are classified as metals, non-metals, and metalloids. • Metals • Metals are elements that are generally shiny when smooth and clean, solid at room temperature, and good conductors of heat and electricity. • Alkali metalsare all the elements in group 1 except hydrogen, and are very reactive. • Alkaline earth metalsare in group 2, and are also highly reactive.

8. The Modern Periodic Table (cont.) • The transition elements are divided into transition metalsand inner transition metals. • The two sets of inner transition metals are called the lanthanide seriesand actinide seriesand are located at the bottom of the periodic table.

9. The Modern Periodic Table (cont.) • Non-metals • Non-metalsare elements that are generally gases or brittle, dull-looking solids, and poor conductors of heat and electricity. • Group 17 is composed of highly reactive elements called halogens. • Group 18 gases are extremely unreactive and commonly called noble gases.

10. The Modern Periodic Table (cont.) • Metalloids • Metalloidshave physical and chemical properties of both metals and non-metals.

11. End of Section 6.1

12. Section 6.2 - Classification of the Elements 3. Organizing the Elements by Electron Configuration • Recall electrons in the highest principal energy level are called valence electrons. • All group 1 elements have one valence electron. • The energy level of an element’s valence electrons indicates the period on the periodic table in which it is found. • The number of valence electrons for elements in groups 13-18 is ten less than their group number.

13. 4. The s-, p-, d-, and f-Block Elements • The shape of the periodic table becomes clear if it is divided into blocks representing the atom’s energy sublevel being filled with valence electrons. • s-block elements consist of groups 1 and 2, and the element helium. • Group 1 elements have a partially filled s orbital with one electron. • Group 2 elements have a completely filled s orbital with two electrons. • After the s-orbital is filled, valence electrons occupy the p-orbital. • Groups 13-18 contain elements with completely or partially filled p orbitals.

14. The s-, p-, d-, and f-Block Elements (cont.) • The d-block contains the transition metals and is the largest block. • There are exceptions, but d-block elements usually have filled outermost s orbital, and filled or partially filled d orbital. • The five d orbitals can hold 10 electrons, so the d-block spans ten groups on the periodic table. • The f-block contains the inner transition metals. • f-block elements have filled or partially filled outermost s orbitals and filled or partially filled 4f and 5f orbitals. • The 7 f orbitals hold 14 electrons, and the inner transition metals span 14 groups.

15. End of Section 6.2

16. Section 6.3 – Periodic Trends 5. Atomic Radius • Many properties of the elements tend to change in a predictable way, known as a trend, as you move across a period or down a group. • Atomic size is a periodic trend defined by how closely an atom lies to a neighboring atom. • The electron cloud does not have a clearly defined edge since it is based on probability. • For metals, atomic radius is half the distance between adjacent nuclei in a crystal of the element. • For elements that occur as molecules (nonmetals), the atomic radius is half the distance between nuclei of identical atoms.

17. Atomic Radius (cont.) • Trends Within Periods • There is a general decrease in atomic radii from left to right across a period. • Caused by increasing positive charge in the nucleus and the fact that the principal energy level remains the same. • Valence electrons are not shielded from the increasing nuclear charge because no additional electrons come between the nucleus and the valence electrons. • Thus, the increased nuclear charge pulls the outermost electrons closer to the nucleus.

18. Atomic Radius (cont.) • Trends Within Groups • Atomic radii generally increases as you move down a group. • The nuclear charge increases as electrons are added to successively higher principal energy levels. • The outermost orbital size increases down a group with energy level, making the atom larger. • Therefore, outer electrons are farther from the nucleus offsetting the increase nuclear charge. • Also, inner electrons shield outer electrons from the pull of the nucleus.

19. 6. Ionic Radius • An ionis an atom or bonded group of atoms with a positive or negative charge. • When atoms lose electrons and form positively charged ions, they always become smaller for two reasons: • The loss of a valence electron can leave an empty outer orbital resulting in a small radius. • Electrostatic repulsion decreases allowing the electrons to be pulled closer to the radius. • When atoms gain electrons, they can become larger, because the addition of an electron increases electrostatic repulsion.

20. Ionic Radius (cont.) • Trends Within Periods • The ionic radii of positive ions generally decrease from left to right. • The ionic radii of negative ions generally decrease from left to right, beginning with group 15 or 16. • Trends Within Groups • In general, there is a gradual increase in ionic size as you move down a group. • Both positive and negative ions increase in size moving down a group.

21. 7. Ionization Energy • Ionization energyis defined as the energy required to remove an electron from a gaseous atom. • The energy required to remove the first electron is called the first ionization energy. • Removing the second electron requires more energy, and is called the second ionization energy. • Each successive ionization requires more energy, but it is not a steady increase.

22. Ionization Energy (cont.) • A high ionization energy value indicates the atom has a strong hold on its electrons, and therefore less likely to form positive ions. • Likewise, a low ionization energy value indicates an atom loses its outer electrons easily, are likely to form positive ions. • Trends Within Periods • First ionization energy generally increases as you move left to right across a period. • Increased nuclear charge produces an increased hold on valence electrons. • Trends Within Groups • First ionization energy generally decreases as you move down the group. • With electrons farther away from the nucleus, less energy required to remove them.

23. Ionization Energy (cont.) • Octet Rule • The octet rule states that atoms tend to gain, lose, or share electrons in order to acquire a full set of eight valence electrons.

24. 8. Electronegativity • The electronegativity of an element indicates the relative ability of its atoms to attract electrons in a chemical bond. • Electronegativity generally decreases as you move down a group, and increases as you move left to right across a period.