Chem Test 2

Steric Number

m + n: Total number of electron groups

AX₂

Linear

Bond Angle: 180

AXE

Linear

Bond Angle: 180

AX₃

Electron Group Arrangement: Trigonal Planar

Molecular Group Arrangement: Trigonal Planar

Bond Angle: 120

AX₂E

Electron Group Arrangement: Trigonal Planar

Molecular Group Arrangement: Bent (V-shaped)

Bond Angle: <120

AX₄

Electron Group: Tetrahedral

Molecular Group: Tetrahedral

Bond Angle: 109.5

AX₃E

Electron Group: Tetrahedral

Molecular Group: Trigonal Pyramidal

Bond Angle: <109.5

AX₂E₂

Electron Group: Tetrahedral

Molecular Group: Bent

Bond Angle: <109.5

AX₅

Electron Group: Trigonal Bipyramidal

Molecular Group: Trigonal Bipyramidal

Bond Angles: 120, 90 (depending on which atoms you are looking between)

AX₄E

Electron Group: Trigonal Bipyramidal

Molecular Group: See-Saw

Bond Angles: <120, <90

AX₃E₂

Electron Group: Trigonal Bipyramidal

Molecular Group: T-Shaped

Bond Angle: <90

AX₂E₃

Electron Group: Trigonal Bipyramidal

Molecular Group: Linear

Bond Angle: 180

AX₆

Electron Group: Octahedral

Molecular Group: Octahedral

Bond Angle: 90

AX₅E

Electron Group: Octahedral

Molecular Group: Square Pyramidal

Bond Angle: <90

AX₄E₂

Electron Group: Octahedral

Molecular Group: Square Planar

Bond Angle: 90

Double Bond Angle

Makes the angle smaller since stronger bond makes atoms closer

Molecular Polarity

An uneven distribution of charge over a whole molecule or a large portion of it

• Homoatomic Atoms (N₂,P₄,S₈,etc) all nonpolar with an exception to O₃)

• Structures with no lone pairs or 1 on central atom are always polar

Hybridization

Mixing atomic orbitals on the same atom

• Predicted through steric number

• Ex: Steric # = 2, Hybridization = sp

• Follow the same format as AX for geometry

• Ex: sp² = Trigonal Planar

Sigma Bonds

Bonds formed by end-to-end overlap of atomic or hybrid orbitals

• s orbitals and hybrid orbitals only form sigma bonds

pi bonds

side-to-side overlap of orbitals, results in regions of -e density above and below the bond axis

• only formed by unhybridized (non combined) p orbitals

Bonding Orbitals

Greater along the bond (between nuclei) + in-phase combination

Energy: Lower than Atomic Orbitals

Antibonding Orbitals

Greater outside the internuclear region + out-of-phase combination

Energy: Higher than atomic orbitals

Hybridized vs. Unhybridized

The amount of unhybridized orbitals depends on the amount of p (px, py, and pz) in the sp structure

Ex: sp has 2 unhybridized p orbitals so it can form 2 pi bonds but sp³ cannot form any as it has 0 unhybridized p oribitals

MO (Molecular Orbital) Diagrams

Bond Order

Bond Types

Magnetic Properties

Diamagnetic

No unpaired electron, weakly repelled by magnetic field

Paramagnetic

one of more unpaired electrons, attracted to a magentic field

Atomic Mass Unit

1amu = Mass of ¹²C / 12

% Abundance

The relative amount of each isotope in a natural sample of that element

Ex: ¹²C: 98.89% + ¹³C: 1.11%

Mass Percent

(mass of element A / mass of compound) * 100

Molecular Formula

Total Number of each atom in a molecule compound

Ex: N₂H₄

Empirical Formula

Lowest whole number ration of atoms in a compound

Ex: NH₂

Steps to find Empirical Formula

1. Convert mass of each element to moles (atomic mass)

2. Convert to smallest ratio by dividing each molar quantity by the smallest one

3. (If necessary) multiply by small number to remove fractions

Solution

Homogenous mixture of two or more substances “single phase” usually involves a liquid.

Solute

Minor component of a solution: What is being dissolved

Solvent

The Major Component: The substance that “does the dissolving”

Strong Electrolytes (Strong Acids) (Soluble)

HCl, HNO₃, H₂SO₄, HClO₄, HBr, HI

Weak Electrolytes (partially soluble)

All other acids and NH₃

Nonelectrolytes (Covalent not an acid or base)

-COOH (weak acid)

-OH (alcohol)

Dilution

Adding pure solvent to a solution (or vice versa) to make a solution of lower concentration.

• Mol of solute stay the same before and after dilution

Precipiation

2 Aqueous solutions are combines and one (or more) insoluble products form

• Seperate your equation and consider the possibilities for ions

Category A: Always Soluble

Cations: Group 1A (Li,Na,Rb,Cs) and NH^4 (ammonium)

Category B: Insoluble except w/ Category A

Cations: Ag⁺, Pb²⁺

Anions: NO₃^-,ClO₄^-,CH₃COO^- | C₂H₃O₂^-(acetate)

Category C: Usually Soluble unless w/ Category B

Halides: Cl^-,Br^-,I^-

Exceptions: Cu⁺ and Hg²⁺ are insoluble

Category D : The In-between

SO₄^2-: Soluble except w/ Category B and Ca²⁺, Sr²⁺, Ba²⁺

OH^-: Insoluble except w/ Category A or Ca²⁺, Sr²⁺, Ba²⁺

Acid

Compound that releases H⁺ when dissolved in water

Formula: H_nX

Base

A compound that releases OH^- in water

Formula M(OH)_n

Strong Bases

LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)₂, Sr(OH)₂, Ba(OH)₂

Weak Acid (Does not completely dissolve)

HF, H₃PO₄, CH₃COOH

Weak Base

NH₃

Oxidation (LEO)

Loss of Electron

Reduction (GER)

Gain of Electron

Redox Reaction

A reaction that involves transfer of electrons

Oxidation Number Rules

Metal in Ionic compound: + O.N. equal to charge

H: 1+ with nonmetal and 1- with metal or metalloid

F: Always -1 in compounds

O: -2 except when bonded to O or F

Cl, Br, I: -1 except when central atom bonded to O or F

Oxidizing Agent

The reactant that gets reduced in a redox reaction

• “does the oxidizing”

• O.N. decreasing during reaction

Reducing Agent

The reactant that gets oxdized in a redox reaction

• “does the reducing”

• O.N. increases during the reaction