Chemistry Unit 3 – Flashcards

Given credit for periodic table (1870). Organized table by increasing atomic mass

Put elements in order of increasing atomic number

The properties of elements repeat every so often - periodically

1. Horizontal rows (7 rows/periods) -has elements with same # of occupied energy levels

ex.)Period 2 has 2 occupied energy levels

2. Vertical columns (18 columns/groups) - has similar characteristics

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Located on the left side of the "staircase" of the Table. Generally forms cations.

Properties:

-good conductors of heat and electricity -can be magnetized

-lustrous (shiny - how much light the metal reflects)

-malleable (hammered or rolled into tin sheets)

-ductile (drawn into wire)

-hardness

Located on the right side of the "staircase" of the Table. Generally forms anions (tends to gain electrons to be closest to noble gases).

Properties:

-good insulators (poor conductors of heat and electricity)

-gases at room temperature (RT): nitrogen, oxygen, fluorine, chlorine, hydrogen

-Liquid at RT: bromine

-Solid at RT: carbon, phosphorus, selenium, sulfur, iodine

Located on the "staircase" between metals and nonmetals.

Properties:

-in-between those of metals and nonmetals -all are solids at RT

-Less malleable than metals but not as brittle as nonmetals

-semi-conducters of electricity (metalloids can be found in calculators, digital watches, etc.)

*Aluminum is a metal

Group 1 except for Hydrogen. Has a +1 charge. Very reactive (especially with water). Rarely found in elemental forms in nature - typically only occur in compounds.

Group 2. Has a +2 charge. Less reactive than alkali metals. Reacts readily with Halogens(group 17). Soft.

Group 17. Has a -1 charge. Only group that contains elements in all 3 states of metals at RT.Found in compounds or as ions in nature.

7 elements exist as diatomic elements (nitrogen, oxygen, fluorine, chlorine, bromine, iodine and hydrogen - all except hydrogen are connected together in the shape of a "7")

Group 18. No charge and un-reactive. Gases at RT. Odorless, colorless, monoatomic. Complete outer shells/valence levels (the outer most energy level [in Bohr's model)

1. 57 - 71

2. 89 - 103

Groups 3-12. Variable charges indicated with roman numerals when in written form.

ex.) Ni(C₂H₃O₂)₂: nickel(II)acetate

An approximation of how electrons behave in their orbitals - where electrons are "probably" found (cloud is more desnse at locations where the probability of finding an electron is high). The maximum number of electrons in an energy level is twice the number of orbitals. Each orbital can only contain 2 electrons

Tells the distribution of electrons in atomic orbitals. A "shorthand" explanation of electron arrangement. Used to determine what energy level an electron is located in ("ground level")

Ex.)Carbon: 1s² 2s² sp²

The first number tells the principle energy level. The letter tells the sublevel. The superscript number tells how many electrons are in that level and when added together overall, tells how many electrons are in the element. For example, Carbone has 6 electrons in 2 sublevels (s ; p)

Determining Electron Configuration

Maximum number of electrons that can fit in each "sublevel": s at most holds 2 electrons, p can hold at most 6 electrons, d at most can hold 10 and f at most can hold 14 electrons.

s=2 (max)
p=6 (max)
d=10 (max)
f= 14 (max)

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IMPORTANT: You must fill orbitals in the order of the chart. To read the chart, follow the arrows from back-end to arrow-head. Remember how many electrons can fit into each orbital. Remember to completely fill the first orbital before moving on to the next orbital.

Number of electrons in the outermost energy level. Can look at electron configuration or group numbers to determine # of valence electrons for element.

Model of an atom in which each dot represents a valence electron. The symbol goes in the middle to represent the nucleus and all other in the atom.

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Hunds rule: every orbital in a sub-shell (s,p,d,f) is singly occupied with 1 electron before any one orbital is doubly occupied - arrows represent electrons which point in opposite directions because electrons tend to spin in opposite direction to balance.

^v  ^v  ^v^v^v  ^v  ^

   1s² 2s²  2p⁶  3s² 3p⁶

these gases derive their name from their tendency not to react - they're stable. Other atoms may achieve a similar stability by either gaining or losing electrons so their electron configuration matches one of the noble gases

The number of electrons atoms gain or lose is equal to the charge that they receive when they become cations or anions in order to become stable like their noble gas. Ex.) Bromine: [Ar]4s² 3d¹⁰ 4p⁵