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Thermochemistry Flashcards

Demonstrate understanding of thermochemical principles and the properties of particles and substances

1
Q

Orbital

A

Region of space where electrons can be found around the nucleus

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2
Q

Subshells

A

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

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3
Q

Electron configuration rules

A
  1. Aufbau - In ground state, electrons fill atomic orbitals from lowest available energy before occupying higher energy levels
  2. Pauli Exclusion Principle - No more than two electrons can occupy the same orbital and they must have opposite spins
  3. Hund’s Rule - Each orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied
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4
Q

Aufbau

A

In ground state, electrons fill atomic orbitals from the lowest available energy before occupying higher energy levels

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5
Q

Pauli Exclusion Principle

A

No more than two electrons can occupy the same orbital and two electrons in the same orbital must have opposite spins

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6
Q

Hund’s Rule

A

Each orbital in a subshell is singly occupied by one electron before any one orbital is doubly occupied

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7
Q

Exceptions to electron configurations

A

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

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8
Q

Noble gases

A

He, Ne, Ar, K, Xe, Rn
All electrons are paired
The outer shell contains 2e (He) or 8e
Extremely stable

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9
Q

Core electrons

A

Electrons in [noble gas]

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10
Q

Valence electrons

A

Electrons outside [noble gas]

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11
Q

Electron configuration of ions

A

+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

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12
Q

Atomic radii

A

Half the distance between two neighbouring atoms

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13
Q

P
L
S
E

A

Proton number - nuclear charge
Level (energy) - distance from the nucleus
Shielding - inner electrons shield outer electrons from nuclear charge
E - electrostatic attraction

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14
Q

Atomic radii across a period

A

Decreases

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15
Q

Atomic radii down a group

A

Increases

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16
Q

Ionic radii +ve cations

A

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

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17
Q

Ionic radii -ve anion

A

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

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18
Q

Ionisation energy

A

The energy require to remove one electron from each atom in one mole of atoms in gaseous state

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19
Q

Ionisation energy across a period

A

Increases

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20
Q

Ionisation energy down a group

A

Decreases

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21
Q

Successive ionisation energies

A

For each successive electron removed, ionisation energy increases due to increased nuclear charge per electron

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22
Q

Electronegativity

A

The ability of an atom to attract a bonding pair of electrons to itself
(electrostatic attraction between the nucleus and the bonding electrons)

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23
Q

Electronegativity across a period

A

Increases as the electrostatic attraction between the nucleus and the bonding electron increases

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24
Q

Electronegativity down a group

A

Decreases as the electrostatic force between the nucleus and the bonding electrons decreases

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25
Octet rule exceptions
H - 2 electrons (1 bond) Be - 4 electrons (2 bonds) B - 6 electrons (3 bonds)
26
Lewis structure steps
1. Count up all valence electrons in the molecule 2. The least electronegative is the central atom, typically the only one of this atom if there are more than two types of atoms 3. Join all other atoms with a single bond 4. Add remaining valence electrons to the outer atoms 5. The central atom should have 4 bonds (8 electrons)
27
VSPER Theory
electron pairs form regions of negative charge which repel each other as far as possible
28
Arrangements/geometry
Areas of electron density 2 - Linear 3 - Trigonal Planar 4 - Tetrahedral 5 - Trigonal bipyramidal 6 - Octahedral
29
Linear shapes
Bonding regions and lone pairs 2 bonding regions 0 lone pairs
30
Linear bond angle
180
31
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
32
Trigonal planar bond angles
120
33
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
34
Tetrahedral bond angle
109.5
35
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
36
Trigonal bipyramidal bond angles
Trigonal bipyramidal: 90 and 120 See-saw: 180, 120, and 90 T - shape: 180 and 90 Linear: 180
37
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
38
Octahedral bonding angles
Octahedral: 90 and 180 Square pyramidal: 90 and 180 Square planar: 90 and 180
39
Dipole
Small charge difference across a bond/molecule that results from a difference in electronegativity of the atoms
40
Intermolecular force
An attractive force between neighbouring molecules
41
Weak intermolecular forces
Temporary dipole-dipole attractions Permanent dipole-dipole attractions Hydrogen bonding
42
Temporary dipole-dipole attractions
Between non-polar molecules or atoms AND Between polar molecules
43
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.
44
Permanent dipole-dipole attraction
Between polar molecules
45
Polarity
1. Polarity of each bond - What are each of the bonds in the molecule polar or non-polar? 2. The shape of the molecule - Are they arranged symmetrically around the central atom?
46
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
47
Weak intermolecular forces in non-polar molecules
Temporary dipole-dipole attraction (TDD)
48
Weak intermolecular forces in polar molecules
Temporary dipole-dipole attraction (TDD) and permanent dipole-dipole attraction (PDD)
49
Hydrogen bonding
Hydrogen directly bonded to NOF Also has TDD and PPD
50
Drawing hydrogen bonding
https://chemistrytalk.org/what-are-hydrogen-bonds/
51
Weak intermolecular force strengths
Strongest -> Weakest Hydrogen bonding Permanent dipole-dipoles (PDD) Temporary dipole-dipoles (TDD)
52
Types of solids
1. Ionic - Ions - Strong Ionic Bonds 2. Metallic - Atoms - Metallic Bonds 3. Covalent network - Atoms - Strong Covalent bonds 4. Molecular - Molecules - Intermolecular Forces
53
Ionic Bond
Strong electrostatic force of attraction between cations and anions
54
Metallic bonds
Attraction between positive nuclei and sea of delocalised valence electrons
55
Covalent bonds
Strong bonds due to shared pair of electrons
56
Intermolecular force
Force of attraction between neighbouring molecules
57
Intramolecular bonds
Bonds with an molecule Ionic and convalent
58
Bonds type and relative strengths
Ionic 1000 Covalent 1000 Hydrogen bonding 50 PDD 10 TDD 1
59
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
60
Temperature
The average kinetic energy of particles in a system
61
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.
62
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
63
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
64
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
65
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.
66
Solubility
The maximum amount of substance that dissolves per litre of solution
67
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.
68
Dissolving and polarity
Polar dissolves in polar Non-polar dissolves in non-polar
69
Enthalpy
The energy in a substance due to kinetic energy of particles and potential energy in chemical bonds. The difference in Hp and Hr
70
Exothermic
Releases heat energy -ve change in enthalpy Forming bonds Products have less energy than reactants
71
Endothermic
Absorbs heat energy +ve change in enthalpy Breaking bonds Products have more energy than reactants
72
ΔrH
The enthalpy change when products form from their constituent reactions under standard conditions
73
ΔfusH
The amount of heat energy required to change one mol of a substance from solid to a liquid
74
ΔvapH
The amount of heat energy required to change one mole of a substance from liquid to gas
75
ΔsubH
The amount of energy required to turn one mol of substance from solid to gas
76
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
77
ΔfH
The enthalpy change when one mole of a substance is formed from its constituent elements under standard conditions Is zero for elements
78
Standard states
Gases: O2, Cl2, H2, noble gases Liquids: Bromine (Br) and Mercury (Hg) Solids: The rest
79
Hess's Law
80
Calorimetry
1. 2. Dissolving = total mass for q = mcΔT 3.
81
Calorimetry accuracy
82
Entropy
The measure of disorder, relation to the degree of dispersal of matter and energy, and the degree of randomness
83
Spontaneity
Determined by the entropy change of the system and surroundings Overall increase in entropy is spontaneous
84
Factors affecting entropy of system
1. Volume 2. Molar mass 3. Temperature change 4. Molecular complexity 5. Phase change 6. Number of moles 7. Solid dissolving in water
85
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
86
Spontaneity ΔH and Δs
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Chemistry (2 decks)

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