The Periodic Guide: Study Guide

Topics: Atom, Periodic table, Chemical bond Pages: 6 (1215 words) Published: May 12, 2014
Lesson 03.01: History of the Periodic Table
Explain how scientific observations led to the development of, and changes to, the periodic table. -Dmitri Mendeleev- first periodic table, organized 63 known elements according to properties, organized into rows and columns and wrote name, mass, and chemical properties on each -Julius Lothar Meyer- independently worked in Germany, similar to Mendeleev -Henry Gwyn Jeffreys Moseley- Worked with Ernest Rutherford, experimented with 38 metals, he found that the positive charge of each element’s nucleus increased by one from element to element as they were arranged in Mendeleev’s periodic table, lead to modern definition of atomic number (# of protons in atom’s nucleus) and the recognition the atomic number was basis for organization of periodic table.

Describe the organization of the modern periodic table.
Arranged from left to right in rows (periods) by increasing atomic number and top to bottom in columns (groups) based on similar chemical properties

Lesson 03.02: Group Names and Properties
Compare and contrast the properties of metals, metalloids, and nonmetals. -Metals- good conductors of heat and electricity and reflect light and heat, most luster (shine) and most are malleable (hammered or rolled into sheets) -Non-metals- poor conductors of heat and electricity, most are gas at room temperature, those that are solid are not malleable -Metalloids- a semiconductor (conduct electricity better than non-metals but not as good as metals), some characteristics of metals but more like nonmetals

Identify groups and sections of the periodic table by group name and common properties. 3.02 notes

Lesson 03.03: Periodic Trends
Describe and explain the trends for effective nuclear charge, atomic radius, ionic radius, and ionization energy across a period and down a group. -Effective Nuclear Charge- the charge (from the nucleus) felt by the valence electrons after you have taken into account the number of shielding electrons that surround the nucleus. -Atomic radius- half the distance between the centers of two atoms of that element that are bonded together -Ionization Energy- the energy required to remove one electron from an element, resulting in a positive ion. -Ionic radius- One-half the diameter of an ion.

A positive ion is called a cation, and a negative ion is called an anion. Nonmetals usually become anions and metals usually become cations.

Predict the properties of an element based on the known patterns of the periodic table. Use periodic table

Describe and explain the periodic trends for electron affinity (honors). Electron affinity-The energy involved when a neutral atom gains an electron Becomes more negative (more energy is given off) for each element across a period from Group 1 to Group 17 because of an increase in effective nuclear charge. Becomes less negative (more positive) going down a group, because each electron is being added to a higher energy level farther from the nucleus.

Explain the exceptions to the trend across a period for ionization energy (honors). Noble gases in Group 18 all have positive electron affinity values. The noble gases must be forced to gain an electron because they already have a full valence energy level. The alkaline earth metals in Group 2 and the nonmetals in Group 15 both have electron affinity values close to zero due to electron repulsion and effective nuclear charge. Nitrogen, in Group 15, does not form a stable -1 ion because when an additional electron is added to nitrogen’s valence energy level, it is added to a 2p orbital that already has one electron. The weak attraction between the added electron and nitrogen’s nucleus is why there is not much energy given off.

Lesson 03.04: Valence Electrons and Bonding
Define and compare ionic and covalent bonding.
-Ionic Bond- A chemical bond that results from electrostatic attraction between positive and negative ions, electrons are given up by one atom and...
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