Chem Exam Two Stuff Generally to Memorize

Indeterminacy

Indefinite future - can only predict probability

Orbitals

Each of the actual or potential patterns of electron density that may be formed in an atom by one or more electrons, and that can be represented as a wave function

Wave Function

Orbitals are determined from mathematical wave functions. They can have positive or negative values, as well as equal 0.

Principal Quantum Number

n

Energy level

Angular Momentum Quantum Number

l

orbital number

Magnetic Quantum Number

ml

Orientation of orbital in an x-y-z plot

Spin Quantum Number

ms

Orientation of electron spin

• spin +1/2

• spin -1/2

Hydrogen Energy Transitions

Hydrogen electrons absorb patterns to move to higher energy levels and emit photos when dropping to lower levels

Radiation

Energy, in the form of waves or particles, emitted by an unstable atom to achieve stability

Rydberg Equation

Equation used to find the energy released or absorbed in an electron transition

S-orbital

Sphere shape

p-orbital

Peanut shape

d-orbital

four lobes shape

Nodes

Where the probability density drops to 0

n-1

More nodes means that there is a higher energy

Phase

Sign of its wave function

• in phase —> same sign

• out of phase —> opposite sign

Atomic Radius

Gets larger from top right to bottom left

Roughly the same size for elements in d-block

Ioniziation energy

Gets larger from bottom left to top right

• Be to B gets smaller

• N to O gets smaller

Increases with the removal of each successive electron

Electron Affinity

Energy associated with the addition of an electron to the valence shell of an atom that is in the gas phase

Generally becomes more negative from left to right

Metallic Character

Gets larger frop top right to bottom left

• Zinc, Cadmium, and Mercury are exceptions because tehy have a full d-orbital which gives them fewer metallic characteristics

Predict Ion Size and Charge

From left to right across a periodic table

• Zeff increases, causing isoelectronic cations to get smaller

• More positively charged —> smaller

Effective Nuclear Charge

Increases from left to right across a periodic table causing isoelectronic anions to get smaller

• More negatively charged —> larger

Effective Nuclear Charge

Zeff = Z - S

z - nuclear charge

s - number of core electrons

Shielding

Increases as you move down a group because more electron shells are added

Blocks valence shell electron attraction due to inner-shell electrons

Penetration Power

An electrons ability to approach the nucleus 

s > p > d > f

Aufbau Principle

Electrons occupy the lowest energy atomic orbitals available

Pauli Exclusion Principle

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

• Electrons in the same orbital must have opposite spins

Hund’s Rule

Electrons will singly occupy all orbitals with parallel spins before pairing up in any orbitals

Paramagnetism

Unpaired electrons

Diamagnetism

All paired electrons

Alkali Metals

1 more electron than previous noble gas, lose 1 electron in reactions and forms a cation with 1+ charge

Alkaline Earth Metals

2 more electrons than previous noble gas, lose 2 electron in reactions, forms a cation with 2+ charge

Transition Metals

d-block

Lose electrons from s and then d orbitals to form cations with various charges

Inner Transition Metals

f-block

Lose electrons from s, d, and/or f orbitals to form cations with various charges

p-block metals

Left side of the metalloids in the periodic table, lose electrons and p orbitals to form cations

Metalloids

In p-block between metals and nonmetals. They can exhibit metallic or nonmetallic behaviors and they can either lose electrons from p and s orbitals to form cations or gain electrons in p-orbitals to form anions

Nonmetals

In the upper right-hand side of the p-block, gain electrons in valence p orbitals, resulting in anions with noble gas electron configurations

Halogens

Nonmetals with 1 fewer electron than the next noble gas

Noble Gases

8 Valence electrons, nonreactive, stable electron configurations

Ionic Bonds

Between metals and nonmetals and electrons are transferred

Covalent Bonds

Between two or more nonmetals and electrons are shared

Single Covalent Bonds

One pair of electrons shared between two atoms

Double Covalent Bonds

Two pairs of electrons shared between two atoms

Triple Covalent Bonds

Three pairs of electrons shared between two atoms and is shorter and stronger than a double bond

Lewis Structures

Valence electrons are represented as dots

Binary Ionic Compounds

Name of cation and base name of anion + ide

Molecular Compounds

The more metallic element goes first

prefix + name of 1st element & prefix + name of 2nd element + ide

Hydrate

Ionic compounds containing a specific number of water molecules associated with each formula unit

Molecular Mass

Mass of an individual molecule or formula unit sum of the masses of the atoms in a single molecule or formula unit

Molar Mass

(Number of atoms in 1st element * atomic mass) + (Number of atoms in 2nd element * atomic mass) 

Mass Percent Composition

Percentage of each element in a compound