AP Chemistry Final Flashcards

Strong Electrons

Dissociates 100% to make ions good conductors of electricity in solution, typically resulting from strong electrolytes such as salts.

Strong Electrolytes (soluble ionic salts)

NaCl, KBr

Strong Acids

HCl, HBr, HI, H₂SO₄, HNO₃, HCLO₄

Strong Bases

LiOH, NaOH, KOH, RbOH, CsOH, Ca(O₄)₂, Sr(OH)₂, Ba(OH)₂

Characteristics of Weak Electrolytes

Few ions in solutions, insoluble compounds, poor conductor

Molarity Formula

M= mole/volume

How does water dissociate molecules?

Partial negative oxygen pulls on an ion, partial positive hydrogen pulls on cation. If the attraction to oxygen/hydrogen is stronger, then the compound is soluble (breaks in water).

How to determine solubility of ionic compound (3-steps)

1. Dissociation

2. Hydration bond of H₂O break

3. Hydration sphere

Oxidizer

LEO (Lose Electrons Oxidization), Oxidizing Agent (increasing)

Best reducing agents ____.

Best reducing agents “lose electrons easily”.

Best oxidizing agents ____.

Best oxidizing agents “gain electrons easily”.

Common Oxidizers

CrO₇ +6e^- → 2Cr³⁺

MnO₄ +5e^- → Mn²⁺

NO₃^- +3e^- → NO

H₂SO₄ +2e^- → SO₂

H₂O₂ +2e^- → 2H₂O

Common Reducers

Cu¹⁺ → Cu²⁺+e^-

Pt^2t → Pt

Balancing Redox in Acid

Balance O by adding H₂), Balance H by adding H⁺

Balancing Redox in Base

Balance O by adding OH, Balance H by adding H₂O

Metals in acid/water (hydro)

X₅ + 2H⁺ → H_2(s) + X²⁺_(aa)

Synthesis Types

Metals + nonmetal → Binary Salt

Metal Oxide + H₂O → Metal + Base

Nonmetal Oxide + H₂O → Acid

NaO_(s) + CO_(g) → Na₂ CO₃

Metallic Oxide + Nonmetallic Oxide → Salt

Decomposition Types

Carbonate → Oxide + CO₂

Chlorates → Chloride + O₂

Hydrogen Peroxide → Water + Oxygen

Salt → Metal + Nonmetal (heat)

Formula for Alkanes

C_nH_2n+2

Isotopes

An atom of an element with same numbers of protons but different mass

Average Atomic Mass

(mass isotopes 1)(% abundance/100) + (mass isotope 2)(% abundance/100) + …

C = ?

3.00 × 10⁸ m/s

1nm = ?

1 × 10^-9m

First Ionization Energy

Energy required to remove the first electron, easy to remove from Group 1, hard for Group 8

Coulomb’s Law

F= (ka₁a₂)/d²

distance from one nucleus to another nucleus, aka ELA

Photoelectron Spectra

high peak = electron in subshell

Paul Exclusion Rule

1L opposite direction

Hund’s Rule

added to own orbital before double

Aufbau Principle

lowest energy first → higher

Zeff

charge felt by specific e^-, left to right (increase)

2nd Ionization Energy

more energy needed to take from a cation

EN (Electronegativity)

how good electron at attracting electron

Ionization from 5A → 6A

removing a paired electron is easier than removing an unpaired one

The Octet Rule

tend to gain, lose, or share e^- to acquire a full octet

Iso electron

same # of electrons

Ionic solids

very strong, require a lot of energy to melt or vaporize these solids (crystal lattice formation)

Lattice energy

energy required to separate the ions in a crystal lattice to individual ions

Melting Point

high ionic attraction, small distance, high charge

Covalent Bonds

one atom shares one or more pairs of e^- so both acquire full octets (2NM)

Nonpolar Covalent Bonds

delta electronegativity = 0 through 0.5

Polar Covalent Bonds

delta electronegativity = 0.5 through 1.9

Ionic

delta electronegativity = 2 through 3.5

Electronegativity trend increases

up and right

Lewis Dot Structure

count valence electrons, choose central atom, connect atoms, distribute remaining electrons

How to find Percent Error

(Experimental Value - Theoretical Value)/Theoretical Value * 100 = % Error

Law of Conservation of Mass

total mass of reactants equals total mass of product

Law of Constant Composition

mass ration of elements in a compound is always the same

VSEPR theory

valence shell electron pair repulsion theory

Electron pairs repel so molecules adjust their shapes

hybridization

mixing of several atomic orbitals to form the same number of equivalent hybrid orbitals

pi bonds

a bond formed when parallel orbitals overlap to share electrons

electron sea model

all metal atoms contribute their valence electrons to form sea of electrons

steric number

number of atoms bonded to central atom + number of lone pairs on atom

brittle

easily broken

Avagadro’s Hypothesis

equal volumes of gases at the same temp share same # of particles

Boyle’s Law

volume increases, more room to move, less collisions, more pressure

Charle’s Law

temp increases, KE increases, to keep # of collisions constant, container must get bigger

Ideal Gas Law

PV = nRT

Partial pressure

(mole fraction %) * (total pressure)

sublimation

change directly from solid to gas without liquid state

solvation

process of surrounding solute particles with solvent to form solution