Semester 2 Chemistry Study Guide

Periodic law

When elements are arranged in order of increasing atomic number.

Rows of the periodic table

Period.

Columns of the periodic table

Group.

Atomic radius change

Decreases as you go up and right in the periodic table.

Electronegativity change

Increases as you go up and right in the periodic table.

Elements in the same group

Same valence electrons.

Elements in the same period

Same number of electron shells.

Metals

Shiny and good conductors of heat, usually solid at room temperature.

Nonmetals

Dull and not good conductors.

Metalloids

In between metals and nonmetals, semiconductors.

Main blocks of the periodic table

SPDF.

London dispersion forces (LDF)

Creates temporary dipole, made by movement of molecules, weakest intermolecular forces.

Dipole-dipole interactions

Exists between neutral polar molecules, effective only when polar molecules are very close together.

Hydrogen bonding

A special type of intermolecular attraction that exists between the hydrogen atom in a polar bond and an unshared electron pair on a nearby electronegative ion or atom.

Dipole

A molecule that has two different charges on different ends.

Partial positive charge

Acquired by the less electronegative atom.

Partial negative charge

Acquired by the more electronegative atom.

Main properties of solids

Definite shape and definite volume.

Kinetic-molecular theory

Explains the properties of liquids and solids.

Intermolecular forces

Forces that attract molecules.

Main properties of liquids

Fill space uniformly and completely, definite volume, much higher densities in comparison to gases, not easily compressed.

Types of solids

Crystalline (set shape) and amorphous.

Changes of state

Six different changes of state occur at the particulate level.

Solution

A homogeneous solution of two or more molecules.

Molecular scale of solution formation

The solute becomes separated and surrounded by the solvent.

Solute

The substance that is dissolved.

Solvent

The substance that dissolves the solute.

solubility

the ability to be dissolved

factors affecting solubility

size, temperature, shaking

miscible liquids

liquids that can be mixed together

immiscible liquids

liquids that cannot be mixed together

solvation

the interaction of a solvent with the dissolved solute

factors affecting solvation

nature of solute and solvent, temperature, pressure

concentration of a solution

measured in moles per liter

saturated solution

has the maximum solute per solvent

unsaturated solution

has less than maximum solute per solvent

reversible reaction

a reaction that can be reversed

chemical equilibrium

the rate of the forward and backward reactions are equal so it looks like it's standing still

equilibrium concentrations

remain constant

Keq

the equilibrium constant, always the same for a given reaction at a given temperature

K > 1

right or product favored reactions, there are more products than reactants

K < 1

left or reactant favored reactions

K = 1

at equilibrium

Le Châtelier's principle

if you have a system at equilibrium and place stress on it, the system will shift to relieve stress

solubility product (K)

sp

Arrhenius acids

H donors

Arrhenius bases

have OH

Brønsted-Lowry acids

H donors

monoprotic acid

has one ionizable hydrogen

polyprotic acid

has more than one ionizable hydrogen

amphoteric substances

can be acid or base

acid-base conjugate pairs

they differ by one hydrogen, the conjugate base will have one less H than the original acid

pH of acid and conjugate acid

the same

pOH of base and conjugate base

the same

reaction of acid or base with water

it becomes the other thing and then you get a reaction with that

neutralization

what happens when an acid is added to a base

strong acid/base

the bigger the Ka/Kb, the stronger the acid/base

water self-ionizes

means that water can dissociate into H+ and OH- ions.

Hydronium ions

H3O+

Hydroxide ions

OH-

K

The equilibrium constant

pH

Negative log of molarity of H+ or H3O+

Titration

Neutralization reaction, used to determine concentration of the unknown acid/base

Oxidation

The loss of electrons

Reduction

The gain of electrons

Oxidation number change

It changes when an atom undergoes oxidation or reduction

Voltaic cell

Also known as a galvanic cell, converts chemical energy into electrical energy by a spontaneous redox reaction

Anode

Where oxidation happens in a voltaic cell

Cathode

Where reduction happens in a voltaic cell

Salt bridge

How electrons flow from the anode to the cathode in a voltaic cell

Standard cell potential

The difference between the cell potential of the anode and the cell potential of the cathode

First law of Thermodynamics

Energy cannot be created or destroyed

Second law of Thermodynamics

Entropy of the universe is always increasing

Third law of Thermodynamics

A perfect crystal at 0 Kelvin has zero entropy; entropy increases with heat

Entropy

Chaos of a system; positional probability (how many places could this possibly be)

Enthalpy

Internal heat of a system

Gibbs Free Energy

Energy that is just lying around

Macromolecules

Four types: Protein, Carbohydrate, Lipid, Nucleic Acid

MV = MV

Used to calculate relationships between molarity and volume

Equilibrium constant expression (Keq)

Write the equilibrium constant expression for a reaction

Ksp

Calculate Ksp

Acid dissociation constant (Ka)

Calculate the acid dissociation constant or base dissociation constant (Kb)

Rules for Oxidation Numbers

1. Hydrogen's oxidation number

   is +1 except for when bonded to metals the

   oxidation number is −1.

2. Fluorine always has an oxidation number of −1.

3. Oxidation number of more electronegative atom in a molecule or complex ion is the same as the charge if that atom were an ion.

4. Oxidation number of oxygen in compounds is always -2, except When bound to fluorine,becomes +2.