The Periodic Table and Periodic Law

1. The Periodic Table and Periodic Law Chapter 6

2. History of the Periodic Table’s Development • Late 1790s: Lavoisier compiled a list of the 23 known elements • 1870: 70 known elements • 1864: John Newlands proposed an organizational scheme for the elements • When elements are arranged by increasing atomic mass their properties repeat every eight element. • Named this periodic property the law of octaves.

3. History of the Periodic Table’s Development • 1869: Meyer and Mendeleev each demonstrated a connection between atomic mass and elemental properties. • Mendeleev published his scheme first and thus received more credit. • Mendeleev arranged elements in columns of increasing atomic mass since the elements had similar properties… this developed into the first periodic table. • Accepted because he predicted the existence and properties of several undiscovered elements and left blank spaces for these (Scandium, Gallium, and Germanium). • 1913: Moseley discovered that atoms of each element contain a unique number of protons in their nuclei (atomic number). • Arranging elements by increasing atomic number correct any errors in Mendeleev’s periodic table. • Resulted in a clear periodic pattern of properties (periodic law).

4. The Modern Periodic Table • Boxes contain elemental name, symbol, atomic number, and atomic mass. • Arranged in columns (groups) and rows (periods) by increasing atomic number. • Groups 1A-8A (groups 1-2 and 13-18) are the main group elements or representative elements. S and p block • Groups 1B-8B (groups 3-12) are the transition elements.

6. The Modern Periodic Table • Three main types of elements: • Metals: shiny, solids at room temperature, good conductors, ductile, malleable. • Alkali Metals (group 1), Alkaline Earth Metals (group 2), transition metals (groups 3-12), and inner transition metals. • Nonmetals: gases or brittle-dull solids, poor conductors. • Halogens (group 17) and Nobles Gases (group 18). • Metalloids (Semimetals): properties of metals and nonmetals.

7. Classification of the Elements • Organizing the Elements by Electron Configuration • Atoms in the same group have similar chemical properties because they have the same number of valence electrons. • The energy level of an element’s valence electrons indicates the period on the periodic table in which it is found. • A representative element’s group number and the number of valence electrons it contains are also related

8. Classification of the Elements • The s-,p-,d-, and f-Block Elements • s-block elements: groups 1 (1 valence electron) and group 2 (2 valence electrons). • p-block elements: groups 13-18 (valence electrons from 3-8). • d-block elements: transition metals (groups 3-12). • f-block elements: inner transition metals.

9. Periodic Trends • Many properties of elements tend to change in a predictable way, known as a trend, as you move across a period or down a group. • Atomic Radius • Ionic Radius • Ionization Energy • Electronegativity

10. Atomic Radius • Atomic size is defined by how closely an atom lies to a neighboring atom. • Atomic radius is defined as half the distance between nuclei of identical atoms that are chemically bonded together. • Atomic Size/Radius decreases as you move left to right across the period… this is because there is an increased nuclear charge (more protons pulling the outer electrons closer to the nucleus). • Atomic Size/Radius increases as you move top to bottom down the group… this is because more energy levels are being added between the nucleus and the outermost energy level (this results in less of a pull on the valence electrons by the nucleus).

11. Ionic Radius • Atoms gain or lose one or more electrons to form ions. • An ion is an atom or group of atoms with a positive or negative charge. • When atoms lose electrons to form a positive charge they become smaller. • This is because the lost electron is a valence electron… often resulting in an empty orbital. • There are also less electrons being pulled by the positive nucleus… which means the electrons are pulled even closer. • When atoms gain electrons to form a negative charge they become larger. • These extra electrons force the outermost electrons further from the nucleus resulting in less of an attraction of the nucleus to the valence electrons. • The trend for the representative elements: • Across the period… decrease. • Down the group… increase.

13. Ionization Energy • Ionization energy is the amount of energy required to remove an electron from an atom. • Energy is needed to overcome the attraction between the positive charge in the nucleus and the negative charge of the electron. • Think of ionization energy as an indication of how strongly an atom’s nucleus holds onto its valence electrons. • High ionization energy = strong hold (*nonmetals) • Low ionization energy = weak hold (*metals) • Second Ionization Energy: energy required to remove a second electron. • Third Ionization Energy… and so on

14. Ionization Energy • First Ionization Energy increases as you move left-to-right across a period because of increased nuclear charge. • First Ionization Energy decreases as you move down the group because the valence electrons are further from the nucleus and easier to remove.

15. Electronegativity • The ability of an atom to attract electrons in a chemical bond. • Fluorine is the most electronegative atom. • Decreases as you move down the group and increases as you move across the period.

16. The Periodic Table and Periodic Law Chapter 6

17. History of the Periodic Table’s Development • Late 1790s: Lavoisier compiled a list of the __________elements • 1870: __________elements • 1864: John Newlands proposed an organizational scheme for the elements • When elements are arranged by __________________their properties repeat every eight element. • Named this periodic property the ____________.

18. History of the Periodic Table’s Development • 1869: ________________each demonstrated a connection between atomic mass and elemental properties. • ___________ published his scheme first and thus received more credit. • Mendeleev arranged elements in columns of increasing atomic mass since the elements had _______________… this developed into the first periodictable. • Accepted because he predicted the existence and properties of several ________________ elements and left blank spaces for these (Scandium, Gallium, and Germanium). • 1913: Moseley discovered that atoms of each element contain a unique number of protons in their nuclei (_______________). • Arranging elements by increasing atomic number correct any errors in Mendeleev’s periodic table. • Resulted in a clear periodic pattern of properties (_______________).

19. The Modern Periodic Table • Boxes contain elemental name, symbol, atomic number, and atomic mass. • Arranged in columns (__________) and rows (_________) by increasing atomic number. • Groups 1A-8A (groups 1-2 and 13-18) are the main group elements or _____________ elements. • Groups 1B-8B (groups 3-12) are the ________ elements.

21. The Modern Periodic Table • Three main types of elements: • _________: shiny, solids at room temperature, good conductors, ductile, malleable. • Alkali Metals (group 1), Alkaline Earth Metals (group 2), transition metals (groups 3-12), and inner transition metals. • _________: gases or brittle-dull solids, poor conductors. • Examples: Halogens (group 17) and Nobles Gases (group 18). • __________ (Semimetals): properties of metals and nonmetals.

22. Classification of the Elements • Organizing the Elements by Electron Configuration • Atoms in the same group have similar chemical properties because they have the same number of ________________. • The energy level of an element’s valence electrons indicates the _______on the periodic table in which it is found. • A representative element’s ________ number and the number of _____ electrons it contains are also related.

23. Classification of the Elements • The s-,p-,d-, and f-Block Elements • _________ elements: groups 1 (1 valence electron) and group 2 (2 valence electrons). • _________ elements: groups 13-18 (valence electrons from 3-8). • _________ elements: transition metals (groups 3-12). • _________ elements: inner transition metals.

24. Periodic Trends • Many properties of elements tend change in a predictable way, known as a trend, as you move across a period or down a group. • _________________ • _________________ • _________________ • _________________

25. Atomic Radius • ____________is defined by how closely an atom lies to a neighboring atom. • ____________is defined as half the distance between nuclei of identical atoms that are chemically bonded together. • Atomic Size/Radius _________ as you move left to right across the period… this is because there is an increased nuclear charge (more protons pulling the outer electrons closer to the nucleus). • Atomic Size/Radius ________ as you move top to bottom down the group… this is because more energy levels are being added between the nucleus and the outermost energy level (this results in less of a pull on the valence electrons by the nucleus).

26. Ionic Radius • Atoms gain or lose one or more electrons to form _____. • An ion is an atom or group of atoms with a positive or negative charge. • When atoms lose electrons to form a positive charge they become _________. • This is because the lost electron is a valence electron… often resulting in an empty orbital. • There are also less electrons being pulled by the positive nucleus… which means the electrons are pulled even closer. • When atoms gain electrons to form a negative charge they become _________. • These extra electrons force the outermost electrons further from the nucleus resulting in less of an attraction of the nucleus to the valence electrons. • The trend for the representative elements: • ______________________… decrease. • ______________________… increase.

28. Ionization Energy • _____________________is the amount of energy required to remove an electron from an atom. • Energy is needed to overcome the attraction between the positive charge in the nucleus and the negative charge of the electron. • Think of ionization energy as an indication of how strongly an atom’s nucleus holds onto its valence electrons. • High ionization energy = strong hold • Low ionization energy = weak hold • _________________________: energy required to remove a second electron. • Third Ionization Energy… and so on

29. Ionization Energy • First Ionization Energy __________ as you move left-to-right across a period because of increased nuclear charge. • First Ionization Energy __________ as you move down the group because the valence electrons are further from the nucleus and easier to remove.

30. Electronegativity • The ability of an atom to attract electrons in a chemical bond. • ________ is the most electronegative atom. • ____________ as you move down the group and _______ as you move across the period.