HL Chem

If the oxidation number increases

Oxidized, reducing agent

If the oxidation number decreases

Reduced, oxidizing agent

All hydrogen atoms are +1 except…

Hydrides are -1

All oxygen atoms are -2 except…

Peroxides are -1

More reactive metals are…

Stronger reducing agents

More reactive non-metals are…

Stronger oxidizing agents

Biochemical Oxygen Demand

The amount of oxygen used by microorganisms to oxidize the organic matter in the water over a fixed period of time and temperature. The higher the BOD, the more organic waste in water.

Voltaic/Galvanic Cells

Spontaneous Reaction, converts chemical energy to electrical energy
Anode is negative, cathode is positive

Electrolytic Cells

Non-spontaneous reaction, converts electrical energy and chemical energy
Anode is positive, cathode is negative

Purpose of salt bridge

Completes the circuit and allows movement of ions to ensure electroneutrality

positive ions go to cathode

Standard Cell Potential (Eºcell)

Energy difference between the cathode and anode

Standard conditions

SATP (25ºC, 101kPa, 1M)

Positive Eº means

More likely to undergo reduction and will be spontaneous

Anode

oxidation, ions are formed, concentration of ions increases as redox reaction occurs

Cathode

electrons are gained, ions are reduced, solid is formed which causes the electrode to increase in mass

Electromagnetic Spectrum

Low energy, long wavelength, low frequency
High energy, short wavelength, high frequency

How does an emission spectrum work

When atoms become excited by heat and electricity, the electrons absorb energy and eventually emit energy as a photon

Convergence limit

As electrons get further from the nucleus, lines will get closer together and will eventually merge into one. Forms a continuous spectrum and gives electron any energy and has no discrete energy.

Heisenberg Uncertainty Principle

1. We cant predict the precise location of an electron or its path

2. We can predict the probability of an electron being in a particular space

Quantum Model of the Atom

Electrons have definite energy levels, electrons have wavelike properties, electrons move in a pulsating wave around the nucleus

Pauli Exclusion Principle

no 2 electrons in an atom have the same set of quantum #s

Aufbau Principle

An energy sublevel must be filled before moving on

Hund’s Rule

1 electron is placed into each orbital before the 2nd

How do electrons spin in the same orbital

They spin in different directions

Electron Configuration Exception: Chromium

4s¹ 3d⁵

Electron configuration Exception: Copper

4s¹ 3d¹⁰

Ionization Energy

Minimum energy required to remove an electron from a neutral gaseous atom

Ionization Energy Formula

X(n-1)(g) → Xn+(g) + e-

Ionization energy group 2 and group 3

Group 3 IE is lower than group 2 because the electron in the 2px orbital is further away from the nucleus than the electrons in the 2s orbitals

Ionization energy group 5 and group 6

Electron repulsion in the 2px orbital of group 6 makes it easier to remove electrons

Formal Charge Formula

(# VE) - ½ (# BE) - (# Non-BE)

Covalent bond

Overlap between 2 atomic orbitals
Sigma bond: end-to-end
Pi bond: Side-to-side

Delocalization

Sharing of a pair of electrons between 3+ ions. Equivalent bonds are intermediate in length and strength (1.5)

Intermolecular Forces

Between the particles of a compound that hold it together in the solid or liquid state

1. Dipole Dipole

2. London Forces (temporary dipoles)

3. Vander Waals Forces

4. Hydrogen Bonding

Dipole Dipole

Between polar molecules, strength depends on polarity

London Forces

Temporary partial charge gives rise to temporary polarity

Strength of london forces depends on

1. # of electrons

   More electrons, higher distance between valence electrons and the nucleus, lower attraction of valence electrons to nucleus. Easier to polarize the cloud and the force is stronger.

2. Size/volume of electron cloud

   Larger electron cloud means less attraction so the cloud is easier to polarize.

3. Shape of molecule
   Larger forces with larger surface area. Branching means smaller surface area.

Hydrogen Bonding

H with N, O, or F

Type of Dipole Dipole

Vander Waal

London + DD + Dipole induced
Dipole Induced: 1 permanent dipole from 1 molecule induces a temporary dipole in the other molecule

Predicting Boiling Point

The stronger the london forces, the higher the BP
More energy is required to overcome the attraction between molecules

Strength of Forces Ranked

London < DD < Hydrogen < Intramolecular Forces

Intramolecular Forces

Bonds present between atoms in a molecule
Ionic, covalent, polar covalent

Giant Covalent Compounds

Large tetrahedral network of carbon or silicon that are covalently bonded to themselves
Theyre interlocking covalent networks, so the strength is amplified

Giant Covalent Compounds Characteristics

Very hard, high MP, insoluble, dont conduct electricity

Allotropes

Different structural modifications of the same element

Diamonds

Allotrope
No plane of weakness through the structure (extremely hard)
Tetrahedral
silicone/silicone dioxide have the same structure

Graphite

Allotrope
Trigonal Planar
layers held together by weak london forces.

Remaining pi electrons are delocalized, the sea of electrons in between layers of graphite allow layers can shift back and force.
Good Conductor

Graphene

Allotrope
1 single layer of graphite
Arranged hexagonally
Excellent conductivity, strength, flexibility, and transparency

C60/Fullerene

60 carbon atoms arranged in hexagonal (20) and pentagonal (12) rings

forms a sphere

composed of covalent bonds but weak london forces (NOT A COVALENT NETWORK)

Metal Characteristics

Lattice structure, valence electrons are loosely held (delocalized), electrostatic in nature, non-directional bonding, good conductors

Melting point of metal increases with

• Increasing ionic charge

• Decreasing ionic radius

• Increased # of delocalized electrons

Alloys

Homogeneous mixtures of 2+ metals or a metal and non metal

Transition metal properties

High mp and density, have 1+ oxidation nuimbers, form colored compounds, catalysts

Paramagnetism

Unpaired electrons attracted by a magnetic field

Diamagnetism

Paired electrons are slightly repelled by magnetic fiels

Ligands

Negative ion or molecule that contain 1+ lone pairs of electrons and form coordinate covalent bonds with complex ions

Haber Process

Molecules dissociate into atom on the catalyst surface, atoms react, molecule leaves catalyst surface as 1

Splitting

D orbitals split into 2 groups in a complex ion. High energy points at ligands lone pair. Caused by repulsion between electrons in the metal ion d-orbitals and the lone pairs of electrons on the ligand

Colour Absorption

Colour related to the amount of splitting. Low splitting takes less energy, which means a lower frequency and higher wavelength is absorbed.

Factors that affect colour

Metal identity: (2+<3+) Ligands pull closer to ion, which means higher repulsion and higher splitting
Oxidation #: higher charges, ligands are closer, greater repulsion, higher splitting
Nature of ligand

Magnetism

Due to unpaired electrons in 3D orbitals/incomplete orbitals.

Higher ionic radius means what (metals)

Lower metallic bonding

Higher bond strength means what (metals)

Higher number of delocalized electrons, smaller atomic radius

Exothermic

Surroundings get hotter, negative value, makes bonds
ex. combustion

Endothermic

Surroundings get colder, positive value, breaks bonds

Bond Enthalpy

Enthalpy change when 1 mole of covalent bonds in a gaseous molecule are broken under standard conditions

∆H Atomization

1 mole of gaseous atom is formed from an element. Always endothermic. (s) → (g)

IE

Electron is removed from an isolated atom in the gaseous state. Always endothermic

EA

an electron is added to an isolated atom. 1st EA is exothermic, second is endothermic

∆H latt (Lattice Enthalpy)

1 mole of an ionic compound is broken apart into its constituent gaseous ions. Always endothermic
Depends on:

• The higher the charge, the higher the mp, the higher the enthalpy

• The smaller the ionic radius, the higher the enthalpy

High charge + low radius = higher charge density

∆H sol

1 mole of solvate is dissolved in excess solvent. Endothermic or Exothermic

∆H hyd

1 mole of a gaseous ion is surrounded by water molecules to form a solution. Always exothermic. Same rules as ∆H latt.

Factors that affect reaction rates

1. the nature of the reactant: varies based on reactivity

2. concentration: higher concentration, more collisions, higher rate

3. Temperature: higher temp, more kinetic energy, more collisions, higher rate

4. Catalysts: changes mechanism and lowers activation energy

5. Surface area: a solid in a powder form increases in surface area which increases the rate

Measuring the rate of reaction

• reactions that produce gas: collect gas and measure its volume/pressure

• reactions involving ions: conductivity can be measured (probe, voltage)

• reactions that change colour: intensity of colour can be measured with a spectrophotometer

Equilibrium: Concentration change

equilibrium shifts in the opposite direction of a concentration increase

Equilibrium: Pressure

Increase pressure → equilibrium shifts to decrease pressure

Q (Reaction Quotient) vs Kc

Q can take any concentration, Kc is only equilibrium concentrations

Rule of 100

divide initial concentration by Kc. >100 assumption to ignore 0 can be made.

Arrhenius Theory

Acids produce H+ when dissolved in water, bases produce OH- when dissolved in water

Brontsed-Lowry Theory

Acids are proton (H+) donors, bases are proton acceptors

Amphoteric Substances

Can act as acids or bases ((H2), HCO3- (baking soda))

Amphiprotic

Can accept or donate a H+ (H2O)

Lewis Theory

Acid is an electron acceptor, base is an electron donour

A conjugate base is

negatively charged

A conjugate acid is

positively charged

Strong acid

reaction goes to completion, conjugate base is weak. will fully dissociate
Ex: HCl, HBr, H2SO4, HNO3, HPO4, HCLO4 REVIEW

Weak Acid

established an equilibrium, conjugate base is strong
Ex: carboxylic acids,HF, H2CO3, H2S, H3BO3

Monoprotic acids

only contain 1 ionizable H (important for ionization energy)

Strong bases

Weak conjugate acids. ex: group 1 or group 2 + OH

Weak Bases

Strong conjugate acids. ex: anything with nitrogen

Salt hydrolysis

when salt reacts with water. does not occur in salts that form neutral solutions.

The larger the Ka

the lower the pka, the stronger the acid

More positive the entropy

the higher the spontaneity, the higher the disorder
increases with temperature and gaseous state molecules

Moving a system to greater stability by:

decreasing enthalpy and increasing entropy

At equilibrium (∆G)

∆G = 0, but G does not = 0

Spontenous at high temps

∆G = + / - , ∆H = +, ∆S = +

Spontaneous at low temps

∆G = + / - , ∆H = -, ∆S = -

Acid Deposition

Occurs when the pH of rain (usually 5.6) is lower than 5

Wet deposition: rain, fog snow

Dry deposition: gases, particles

Sources of acidic pollutants

SO2 (sulfur dioxide): power stations, volcanoes

NOx (nitrogen oxides): transport, lightning

Affect of Acidic rain on vegetation

• displaces metal ions (Ca, Mg, K) from the soil and washes them away. Mg ions are needed to produce chlorophyll, affects photosynthesis.

• can cause Al ions to dissolve from rocks → damages plant roots and limits water uptake