Chemistry Lecture Notes Chapters 4, 6, 7, 8, 9, 11, 23

Flashcards from lecture notes

Solute

The substance being dissolved.

Solvent

The dissolving medium in a solution.

Unit of Concentration

Indicates the quantity of solute in a solvent/solution.

Molarity (M)

M=moles solute/L solution

Electrolytes

Substances that form ions in solution and conduct electricity.

Hydration

Dissolving process when the solvent is water

Water

A polar molecule and often referred to as the universal solvent

Dilution Equation

M₁V₁ = M₂V₂

Saturated Solution

When the rate of the solution process equals the rate of crystallization

Supersaturated

A solution containing more solute than a saturated solution

Solubility

Amount of solute needed to form a saturated solution in a given quantity of solvent.

Metathesis Reactions

Reactions involving the exchange of ions in aqueous solution.

Spectator Ions

Ions that do not participate in the actual chemistry of a reaction.

Acids

Substances that ionize in water to form hydrogen ions.

Bases

Substances that react with hydrogen ions.

Neutralization Reactions

Acid + Base --> Water + Salt

Gas Forming Neutralization reactions

Bases that contain carbonate or bicarbonate ions and produce carbon dioxide gas.

Titration

Lab procedure used to determine the concentration of a solution using a standard solution.

Standard Solution

Substance of known concentration used in a titration.

Equivalence Point (End Point)

The point in a titration where stoichiometric equivalence is reached between the reactants.

Indicators

Substances that change color to indicate the equivalence point in a titration.

Oxidation

Loss of electrons

Reduction

Gain of electrons

Reducing Agent

The substance that gives up electrons and contains atoms that are oxidized.

Oxidizing Agent

The substance that gains electrons and contains atoms that are reduced.

Arrhenius acid

Increase [H+] in water

Arrhenius base

Increase [OH-] in water

Energy

The capacity to do work or transfer heat

Internal Energy

All p.e. + all k.e. of a system

Open System

Transfer of both energy and matter to and from surroundings

Closed System

Energy transfer but not matter

Isolated System

No energy or matter transfer

Specific Heat (C)

Amount of heat required to raise the temperature of 1.0g substance 1ºC

Enthalpy (H)

Heat content

Endothermic Reaction

Reaction which requires energy… ΔH is positive

Exothermic Reaction

Reaction releases energy… ΔH is negative.

Calorimetry

Measurement of heat change for chemical and physical processes

Calorimeter

Device used to measure heat gained or lost in a chem./phys process

Enthalpy of vaporization

△H vap

Enthalpy of fusion

△H for solid to liquid

Enthalpy of combustion

△H for a substance reacting with oxygen

Enthalpy of formation

△Hf : heat change related to the formation of a compound from its constituent elements

Hess’s Law

if a reaction is carried out in a series of steps, ∆H for the overall reaction will be equal to the sum of enthalpy changes for the individual steps

Fuel value

Energy released when 1 g of a material is combusted.

Spontaneous reactions

Naturally favor the formation of products at specified conditions

Non-spontaneous reactions

Reactions that do not favor the formation of products at specified conditions

Entropy

Measure of the disorder of a system

Law of Disorder

things move in the direction of maximum disorder or randomness

Gibb’s Free Energy Equation

∆G=∆H - T∆S

C

C = speed of light=3.0 x 108 m/s

λ

λ = wavelength (m)

ν

ν= frequency (cycles per second, s-1, Hz)

photon, quantum

packets of electromagnetic radiation

Quantized energy

E= hν

Planck's constant

h= Planck’s constant = 6.63 x 10-34 joule-sec

Plum pudding model

JJ Thomson’s model of the atom which has electrons dispersed in a cloud of positive charge

Nuclear Atom

Model of the atom with a dense, positively charged nucleus with protons and neutrons

Planetary Model

Model of the atom with electrons on fixed paths of specific energy

Schrödinger model

Quantum Mechanical model

Energy Levels

Electrons are arranged in Energy Levels designated by n=row number

Sublevels

Correspond to the cloud shapes (known as atomic orbitals)

Orbitals

Hold up to 2 electrons max

Aufbau Principle

electrons enter orbitals of lowest energy first

Pauli Exclusion Principle

2 e- at most may occupy an orbital; 2 e- of same orbital must have opposite spin

Hund’s Rule

when electrons occupy orbitals of equal energy, one e- enters each orbital until all orbitals contain 1 e- with parallel spins

Orbital Diagrams

Visual representations of electron arrangements

Valence electrons

The outer most orbitals

Octet Rule

For stability, most elements want full s and p sublevels

Periodicity

Repetition of patterns of physical and chemical properties are used to place elements in the Periodic Table of Elements.

Atomic radius

½ distance between nuclei of two like atoms bonded

Ionization energy

the quantity of energy required to remove an electron from ground state of an atom

Electron affinity

Energy change that measures the attraction of the atom for electrons (forming anions)

Group 1A Alkali Metals

Very reactive--only found in compounds. React vigorously with water

Group 7A Halogens

Diatomic

Properties and positions on the PT

Metals/Non-metals/Metalloids

Mendeleev 1869

Used atomic weight and saved spaces for undiscovered elements

Electronegativity

the tendency of an atom in a molecule to attract shared electrons to itself

Ionic Bonding

metal + nonmetal or polyatomic ions

Covalent Bonding

sharing of electrons

Octet Rule (Gilbert Lewis)

Atoms tend to gain, lose or share electrons until they are surrounded by 8 valence electrons (Noble Gas configuration)

Lewis Dot Structures

representation of valence electrons

VSEPR Model

3D shape of molecules based on Lewis Structures.

Non-polar covalent bond

Electrons shared between two atoms of the same element

Polar covalent bond

Electrons between two different elements not shared equally

Dipole moment, μ

Polar molecule with a quantitative measure of the magnitude of a dipole is the this

Intramolecular

True bonding. Inside/Within the molecule. VERY STRONG

Intermolecular Forces (IMFs)

Not a true bond. Between molecules. Weaker than Intra-

Intermolecular Forces (IMFs)

attractive forces BETWEEN molecules

Dipole-Dipole Interactions

Electrostatic attraction between polar molecules

Electron Sea Model

Cations immersed in sea of delocalized valence electrons

Ideal gas

Gas that obeys all 5 postulates of kinetic molecular theory

Real gas

Gas with particles that have finite volume and presence of intermolecular attractions

Gas Laws

Behavior of gases is predictable and measurable

Boyle’s Law

The volume of a sample of gas in a flexible container decreases if the external pressure increases (temperature and amount constant)

Avogadro’s Hypothesis (aka: Molar Volume)

Equal volumes of gases contain equal numbers of molecules

Dynamic Equilibrium

For reversible reactions: when opposing rxns proceed at equal rates