Becca Chem Chapter 8

Front 1

Definition: Ground state

Back 1

The lowest energy state

Front 2

Written format for electron configuration

Back 2

Atomic symbol, then the list of orbitals from lowest to highest n, with a superscript for the number of electrons in that orbital.

Ex: O 1s^2 2s^2 2p^4

Front 3

Definition: Noble gas notation

Back 3

Replaces the part of an atom's electron configuration that is the same as its previous noble gas with the symbol of that gas in brackets. Ex: Na's notation is [Ne] 3s^1

Front 4

Definition: Orbital diagram

Back 4

A diagram that shows an element's electron configuration and symbolizes the electrons as arrows in boxes representing orbitals

Front 5

The Pauli exclusion principle

Back 5

No two electrons in an atom can have the same four quantum numbers

Front 6

How many electrons can occupy a single orbital?

Back 6

Two, as long as they have opposite spins. Half arrow up for up and half arrow down for down

Front 7

Sublevels by least to greatest energy

Back 7

E(s orbital) < E(p orbital) < E(d orbital) < E(f orbital)

Front 8

Definition: Aufbau principle

Back 8

The pattern of orbital filling based on starting with the orbitals with the least energy and moving gradually to those with greater energy

Front 9

Definition: Hund's Rule

Back 9

When filling degenerate orbitals, electrons fill them singly first, with parallel spins, before doubling up.

Front 10

Chart for filling electron configurations

Back 10

Exceptions: Chromium (4s^1, 3d^5)

and Copper (4s^1, 3d^10)

Front 11

In the periodic table, the row number is equal to which quantum number?

Back 11

The highest principal quantum number (n) of the atom

Front 12

Definition: Valence electrons

Back 12

For main-group elements: Electrons that are in the outermost principal energy level

For transition metals: Electrons in the outermost principal energy level as well as the outermost d sublevel

Front 13

Definition: Core electrons

Back 13

Non-valence electrons, i.e. electrons that are in complete principal energy levels and complete d and f sublevels

Front 14

Electron configurations for monoatomic ions

Back 14

For cations: Subtract electrons equal to the charge number from the atom's electron configuration.
For anions: Add electrons equal to the charge number to the atom's electron configuration. For main group use reverse order of filling. For transition metals always take from the highest n level first.

Front 15

Definition: Isoelectronic

Back 15

Having the same electronic configuration

Front 16

When atoms react they exchange electrons to become isoelectronic with _____

Back 16

Their nearest noble gas

Front 17

Definition: Periodic property

Back 17

A property of an atom that is predictable by its place in the periodic table

Front 18

Three periodic properties to know

Back 18

Atomic size/radius, ionization energy, and electronegativity

Front 19

Periodic trend of atomic size/radius

Back 19

Across a period size decreases, down a group size increases

Front 20

Definition: Ionization energy

Back 20

The energy required to remove an electron from an atom in its gaseous state

Front 21

Periodic trend of ionization energy

Back 21

Across a period energy increases (more negative), down a group energy decreases (less negative)

Front 22

Definition: Electronegativity

Back 22

The attraction an atom has for electrons in a covalent bond

Front 23

Periodic trend of electronegativity

Back 23

Across a period increases, down a group decreases

Front 24

Label blocks of the Periodic table

Back 24