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 gained by another atom, and then those atoms are attracted to each other. Between a metal and nonmetal. -Covalent bond- Electrons are shared between two atoms, neither atom completely gains or loses electrons. Between two nonmetals.
Relate your knowledge of the periodic trends to the chemical bonding exhibited by various elements.
Lesson 03.05: Ionic Bonding and Writing Formulas
Determine an element’s ionic charge based on its location on the periodic table. Group 1- 1+
Group 2- 2+
Group 3- 3+
Group 4- 4+
Group 5- 3-
Group 6- 2-
Group 7- 1-
Group 8- non-reactive noble gases
Write the correct ionic formula when given two elements that bond ionically. Use question above and periodic table
Lesson 03.06: Covalent Bonding and Lewis Structures
Determine how many covalent bonds an atom needs in order to fill its valence shell, using the periodic table. Must get to 8 valance electrons.
Ex. Group 17 needs one more valance electron; group 6 needs 2 more valance electrons
Draw correct Lewis structures to model covalently bonded molecules when given the name or formula of the molecule.
Describe your observations and conclusions from the virtual lab.
Lesson 03.07: Intermolecular Forces
Use VSEPR theory to predict the shape of a molecule based on its Lewis structure. The VSEPR theory is about geometry of compounds and electron location.
Compare and contrast intermolecular forces (London dispersion, dipole-dipole, hydrogen bonding, and ion-dipole).
London dispersion forces occur between all molecules and particles but are the only force of attraction between nonpolar molecules or noble gas atoms. These forces are the weakest of the intermolecular forces. The London dispersion forces are caused by the motion of electrons. Dipole-dipole forces are electrostatic interactions of permanent dipoles in polar molecules.
The attractive forces that occur between the positive end of one polar molecule and the negative end of another polar molecule tend to align the molecules to increase the attraction. Hydrogen bonding is a particularly strong dipole-dipole interaction in which hydrogen is covalently bonded to a highly electronegative element, and attracted to the very electronegative element in another molecule. It occurs only in molecules containing N-H, O-H or F-H bonds. Ion-dipole forces are attractive forces that result from the electrostatic attraction between an ionic compound and a polar molecule. This interaction is most commonly found in solutions, especially in solutions of ionic compounds in polar solvents, such as water.
Identify the intermolecular forces experienced by different compounds. Intramolecular Forces: The forces of attraction that occur between individual molecules.
Lesson 03.08: Naming Compounds
Correctly name covalent compounds, ionic compounds, and acids when given their formulas. A metal forms a positive ion (cation) and a nonmetal forms a negative ion (anion). The cation and anion combine to form an ionic compound, more specifically referred to as a binary ionic compound.
Write the formulas for ionic compounds, covalent compounds, and acids from their names.
Name hydrates or write the formula of a hydrate when given its name (honors). Same prefixes
Lesson 03.09: Molar Mass of Compounds
Calculate the molar mass of compounds from the formula.
Determine empirical formulas from percent by mass or mass data.
Empirical formula: The formula of a compound in which the subscripts represent the lowest whole-number ratio of the atoms.
Determine the molecular formula from the empirical formula and molar mass of a substance. No clue
Calculate the molar mass of a hydrate and determine the formula of a hydrate from experimental data (honors). H2O- 18.015
Determine the empirical formula of a compound from the mass of the products produced in experimental reactions (honors).