Thermochemistry Flashcards
Demonstrate understanding of thermochemical principles and the properties of particles and substances
Orbital
Region of space where electrons can be found around the nucleus
Subshells
s - 1 orbital with 2 electrons each
p - 3 orbitals with 2 electrons each
d - 5 orbitals with 2 electrons each
f - 7 orbitals with 2 electrons each
Electron configuration rules
- Aufbau - In ground state, electrons fill atomic orbitals from lowest available energy before occupying higher energy levels
- Pauli Exclusion Principle - No more than two electrons can occupy the same orbital and they must have opposite spins
- Hund’s Rule - Each orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied
Aufbau
In ground state, electrons fill atomic orbitals from the lowest available energy before occupying higher energy levels
Pauli Exclusion Principle
No more than two electrons can occupy the same orbital and two electrons in the same orbital must have opposite spins
Hund’s Rule
Each orbital in a subshell is singly occupied by one electron before any one orbital is doubly occupied
Exceptions to electron configurations
Cr and Cu
The 4s orbitals only contain one electron as it is a lower energy arrangement for the 3d orbital to to be half-full or completely full
Noble gases
He, Ne, Ar, K, Xe, Rn
All electrons are paired
The outer shell contains 2e (He) or 8e
Extremely stable
Core electrons
Electrons in [noble gas]
Valence electrons
Electrons outside [noble gas]
Electron configuration of ions
+ve cations form when electrons are removed from the highest energy level orbits
-ve anions form when electrons are added to the highest energy level orbitals
EXCEPTION: 4s subshells are filled and emptied before the 3d
Atomic radii
Half the distance between two neighbouring atoms
P
L
S
E
Proton number - nuclear charge
Level (energy) - distance from the nucleus
Shielding - inner electrons shield outer electrons from nuclear charge
E - electrostatic attraction
Atomic radii across a period
Decreases
Atomic radii down a group
Increases
Ionic radii +ve cations
P - same number of protons, smae nuclear charge
L - valence e are removed, often all e from the outermost energy level. Decreases distance
S - decreased electron shielding
E - greater electrostatic attraction
Smaller radius
Ionic radii -ve anion
P - same number of protons, same nuclear charge
L - electrons added to the same electron level, increased e-e repulsion, increase distance from nucleus
S - same electron shielding
E - electrostatic force decreases
Larger radius
Ionisation energy
The energy require to remove one electron from each atom in one mole of atoms in gaseous state
Ionisation energy across a period
Increases
Ionisation energy down a group
Decreases
Successive ionisation energies
For each successive electron removed, ionisation energy increases due to increased nuclear charge per electron
Electronegativity
The ability of an atom to attract a bonding pair of electrons to itself
(electrostatic attraction between the nucleus and the bonding electrons)
Electronegativity across a period
Increases as the electrostatic attraction between the nucleus and the bonding electron increases
Electronegativity down a group
Decreases as the electrostatic force between the nucleus and the bonding electrons decreases
Octet rule exceptions
H - 2 electrons (1 bond)
Be - 4 electrons (2 bonds)
B - 6 electrons (3 bonds)
Lewis structure steps
- Count up all valence electrons in the molecule
- The least electronegative is the central atom, typically the only one of this atom if there are more than two types of atoms
- Join all other atoms with a single bond
- Add remaining valence electrons to the outer atoms
- The central atom should have 4 bonds (8 electrons)
VSPER Theory
electron pairs form regions of negative charge which repel each other as far as possible
Arrangements/geometry
Areas of electron density
2 - Linear
3 - Trigonal Planar
4 - Tetrahedral
5 - Trigonal bipyramidal
6 - Octahedral
Linear shapes
Bonding regions and lone pairs
2 bonding regions
0 lone pairs
Linear bond angle
180
Trigonal planar shapes
Bonding regions and lone pairs
Trigonal planar:
- 3 bonding regions and 0 lone pairs
Bent/v shape:
- 2 bonding regions and 1 lone pair
Trigonal planar bond angles
120
Tetrahedral shapes
Bonding regions and lone pairs
Tetrahedral:
- 4 bonding regions and 0 lone pairs
Trigonal pyramidal
- 3 bonding regions and 1 lone pair
Bent/v-shape
- 2 bonding regions and 2 lone pairs
Tetrahedral bond angle
109.5
Trigonal bipyramidal shapes
Bonding regions and lone pairs
Trigonal bipyramidal:
- 5 bonding regions and 0 lone pairs
See-saw:
- 4 bonding regions and 1 lone pair
T - shape:
- 3 bonding regions and 2 lone pairs
Linear:
- 2 bonding regions and 3 lone pairs
Trigonal bipyramidal bond angles
Trigonal bipyramidal:
90 and 120
See-saw:
180, 120, and 90
T - shape:
180 and 90
Linear:
180
Octahedral shapes
Bonding regions and lone pairs
Octahedral:
- 6 bonding regions and 0 lone pairs
Square pyramidal:
- 5 bonding regions and 1 lone pair
Square planar:
- 4 bonding regions and 2 lone pairs
Octahedral bonding angles
Octahedral:
90 and 180
Square pyramidal:
90 and 180
Square planar:
90 and 180
Dipole
Small charge difference across a bond/molecule that results from a difference in electronegativity of the atoms
Intermolecular force
An attractive force between neighbouring molecules
Weak intermolecular forces
Temporary dipole-dipole attractions
Permanent dipole-dipole attractions
Hydrogen bonding
Temporary dipole-dipole attractions
Between non-polar molecules or atoms
AND
Between polar molecules
Factors affecting strength of TDD
Size: Larger molar mass = larger electron cloud. The molecule is more polarizable and TDD form more easily.
Shape: Linear molecules can get closer to each other than branched chains. Greater attraction.
Permanent dipole-dipole attraction
Between polar molecules
Polarity
- Polarity of each bond
- What are each of the bonds in the molecule polar or non-polar? - The shape of the molecule
- Are they arranged symmetrically around the central atom?
Bond polarity
Bond polarity results when there is a difference in electronegativity between two atoms in a covalent bond. Creates dipole due to charge separation
Weak intermolecular forces in non-polar molecules
Temporary dipole-dipole attraction (TDD)
Weak intermolecular forces in polar molecules
Temporary dipole-dipole attraction (TDD) and permanent dipole-dipole attraction (PDD)
Hydrogen bonding
Hydrogen directly bonded to NOF
Also has TDD and PPD
Drawing hydrogen bonding
https://chemistrytalk.org/what-are-hydrogen-bonds/
Weak intermolecular force strengths
Strongest -> Weakest
Hydrogen bonding
Permanent dipole-dipoles (PDD)
Temporary dipole-dipoles (TDD)
Types of solids
- Ionic - Ions - Strong Ionic Bonds
- Metallic - Atoms - Metallic Bonds
- Covalent network - Atoms - Strong Covalent bonds
- Molecular - Molecules - Intermolecular Forces
Ionic Bond
Strong electrostatic force of attraction between cations and anions
Metallic bonds
Attraction between positive nuclei and sea of delocalised valence electrons
Covalent bonds
Strong bonds due to shared pair of electrons
Intermolecular force
Force of attraction between neighbouring molecules
Intramolecular bonds
Bonds with an molecule
Ionic and convalent
Bonds type and relative strengths
Ionic 1000
Covalent 1000
Hydrogen bonding 50
PDD 10
TDD 1
Atoms in hydrogen bonding
Period 2
Small atomic radii
Highly electronegative
Have lone pairs of electrons which can attract the partially charged hydrogen of another molecule
Temperature
The average kinetic energy of particles in a system
Temperature vs Heat
Temperature relates to the speed of the atoms and molecules in a substance, whereas, heat is the total energy of these atoms and molecules.
Higher temperature objects have more heat energy than lower temperatures.
Heat can be transferred by conduction, convection, and radiation.
Melting point
The amount of heat energy required to turn a substance from solid to liquid
Fusion
Some forces are overcome and particles break away from regular arrangement
Boiling point
The amount of heat energy required to turn a substance from liquid to gas
Vapourisation
All the forces between particles are completely disrupted, breaking them out of arrangement
Latent heat
The heat required the change the state of a substance without causing a change in temperature
https://www.sciencelearn.org.nz/images/231-latent-heat-graph
Graph explaination
When temperature is increasing, the molecules of water are gaining kinetic energy.
When the temperature is constant, the energy is supplied is used to overcome the intermolecular forces rather then increase the kinetic energy of the particles.
ΔfusH takes less energy than ΔvapH as only some of the intermolecular forces are overcome when solid turns to liquid. Whereas for a liquid to turn to gas, more energy is required to overcome all of the intermolecular forces.
Solubility
The maximum amount of substance that dissolves per litre of solution
Attractions in solubility
Solute-solute
Solvent-sovlent
Solvent-solute
For a substand to dissolve, the solvent-solute attraction must be greater than the solvent-solvent and solute-solute attraction.
Dissolving and polarity
Polar dissolves in polar
Non-polar dissolves in non-polar
Enthalpy
The energy in a substance due to kinetic energy of particles and potential energy in chemical bonds.
The difference in Hp and Hr
Exothermic
Releases heat energy
-ve change in enthalpy
Forming bonds
Products have less energy than reactants
Endothermic
Absorbs heat energy
+ve change in enthalpy
Breaking bonds
Products have more energy than reactants
ΔrH
The enthalpy change when products form from their constituent reactions under standard conditions
ΔfusH
The amount of heat energy required to change one mol of a substance from solid to a liquid
ΔvapH
The amount of heat energy required to change one mole of a substance from liquid to gas
ΔsubH
The amount of energy required to turn one mol of substance from solid to gas
ΔcH
The enthalpy change when one mole of an element or compound reacts completely with oxygen under standard conditions
Forms H2O and CO2
always -ve
ΔfH
The enthalpy change when one mole of a substance is formed from its constituent elements under standard conditions
Is zero for elements
Standard states
Gases: O2, Cl2, H2, noble gases
Liquids: Bromine (Br) and Mercury (Hg)
Solids: The rest
Hess’s Law
Calorimetry
- Dissolving = total mass for q = mcΔT
3.
Calorimetry accuracy
Entropy
The measure of disorder, relation to the degree of dispersal of matter and energy, and the degree of randomness
Spontaneity
Determined by the entropy change of the system and surroundings
Overall increase in entropy is spontaneous
Factors affecting entropy of system
- Volume
- Molar mass
- Temperature change
- Molecular complexity
- Phase change
- Number of moles
- Solid dissolving in water
Factors affecting entropy of surroundings
Exothermic:
Releases energy from system to surroundings. Increased dispersal of energy and matter. Increased entropy.
Endothermic:
System absorbs energy from the surroundings. Decreased dispersal of matter and energy. Decreases entropy
Spontaneity ΔH and Δs
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