Chapter 5
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29 terms
Terms | Definitions |
|---|---|
 Cannizzaro | Created a method for accurately measuring relative atomic mass. |
| Mendeleev | Organized the elements according to properties. He placed the name of each element on one side of a card, and the properties on another, then sorted them. He noticed that elements had similar properties in regular intervals when grouped by atomic mass, however there were a few exceptions. He published the first periodic table. |
| Periodicy | The repeating pattern in grouping of elements, such as when Mendeleev grouped them by atomic mass. |
| Mosley | Observed that elements fit into patterns of properties better when arranged by nuclear charge (atomic number). |
| Periodic Law | When elements are arranged in order of increasing atomic number, elements with similar properties appear in regular intervals. |
| Periodic Table | An arrangement of elements in order of atomic numbers so that elements with similar properties fall in the same column (group). |
| Ramsey/Strutt | Discovered the noble gas Argon. Showed that helium exists on Earth. Proposed Group 18 for noble gasses. |
| Lanthanides | 14 elements from atomic number 48-71 with extremely similar properties. They are transition metals located on the top row of the bottom section of the periodic table. |
| Actinides | 14 elements from atomic numbers 90-103 with similar properties. They are transition metals on the bottom row beneath the periodic table. Radioactive. |
| Periods | 7 horizontal rows on the periodic table that are organized based on groups. The length is determined by the number of electrons that can fill the sublevels of that period. |
| S-Block | Groups 1 and 2. Chemically reactive metals. Include alkali metals and alkaline-earth metals. |
| Alkali Metals | Group 1 elements. Silvery appearance. Soft enough to be cut with a knife. Very reactive, thus not found purely in nature. Combine with nonmetals and water. |
| Alkaline Earth Metals | Group 2 elements. Harder, denser, and stronger than alkali metals. Have higher melting points than alkalis. Very reactive. |
| D-Block | Sublevel where there are three possible sublevels. The 4s sublevel is filled before 3d. 5 orbitals for each sublevel, each of which can hold 2 electrons. Transition metals. |
| Transition Metals | D-Block elements. Good at conducting electricity. High luster. Less reactive than alkali metals and alkaline-earth metals. |
| P-Block | Groups 13-18 except for Helium. Includes nonmetals, metals, and metalloids. Number of valence electtrons can be found by taking the group number minus 10. Metals in the p-block are harder and softer than d-block transition metals. Denser than alkaline-earth metals. |
| Main-Group Elements | The p and s-block elements. |
| F-Block | Elements between groups 3 and 4. Includes actinides and lanthranides. |
| Atomic Radius | ½ the distance between the nuclei of identical bonded atoms. Measured this way because the edges of atoms are uneven. Gets smaller across a period due to more positive charge. Increases down a group |
| Ionization Energy | A measure of the energy required to remove one electron from a neutral atom and form an ion. Measured only with isolated atoms in the gas phase. Measured in KJ/mole. Energy is required for the release. |
| Trends of Ionization Energy | In general, increases across a period due to greater nuclear charge, and thus more strongly attracted electrons. Generally decreases down groups. As you move down, the outer electron shells are in higher energy levels, and thus are further from the nucleus. They are harder to pull apart. Also, as atomic number increases, more electrons are between the nucleus and outer shell, which shields outer electrons from the pull of the nucleus. |
| _____ Ionization Energy | The energy required to remove a _____ electron from a positive ion. Multiple electrons can be removed before a positive ion reaches a neutral state. Energy required gets higher with every new electron because after one electron is moved, the others feel a stronger positive nuclear charge. Biggest jump of ionization energy from noble gas-like state to one less electron. |
| Electron Affinity | The energy change that occurs when an electron is acquired by a neutral atom. Often causes a release in energy. Other times, energy can "force" an atom to gain an electron. Positive electron affinities are extremely difficult to measure because it will be unstable and lose the electron almost immediately. |
| Trends of Electron Affinity | In general, affinity becomes more negative across a period in the p-block. Electron affinity generally is greater as you move down groups. Second electron affinities are all positive. Much more difficult to add more and more electrons. |
| Cation | A positive ion. Formed by the loss of electrons. Causes a decreased atomic radius. Metals on the left of the periodic table are cations. |
| Anion | A negative ion. Formed by the addition of electrons. Causes decreased atomic radius. Nonmetals in the upper right of the periodic table tend to form anions. |
| Electronegativity | A measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound. Valence electrons tend to be unevenly concentrated in charge in compounds, leaning toward one atom. Fluorine is the most electromagnetic element, and has a value of 4. The electronegativity of all elements is based on this 4.0 scale of Fluorine. |
| Trends of Electronegativity | Tends to increase across periods, and decrease or stay the same down groups. |
| Periodic Properties of D-Block Elements | Atomic radii tends to decrease slightly across periods. Ionization energy increases across a period. Ionization energy decreases down groups because valence elecytrons are les shielded by the charge of electrons in the interior sublevels. They lose electrons from the highest energy level, even if the d sublevel fills first. EX: Fe ([Ar]3d64s21) loses the 4s electrons then 3d ones. Electronegativity increases as radii decreases. |
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