HL Chem Terms Sem 1

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

Measure of 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

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 fives 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 that 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