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Chemistry > Thermochemistry > Flashcards

Thermochemistry Flashcards

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

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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
Q

Octet rule exceptions

A

H - 2 electrons (1 bond)
Be - 4 electrons (2 bonds)
B - 6 electrons (3 bonds)

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26
Q

Lewis structure steps

A
  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)
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27
Q

VSPER Theory

A

electron pairs form regions of negative charge which repel each other as far as possible

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28
Q

Arrangements/geometry

A

Areas of electron density
2 - Linear
3 - Trigonal Planar
4 - Tetrahedral
5 - Trigonal bipyramidal
6 - Octahedral

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29
Q

Linear shapes

A

Bonding regions and lone pairs
2 bonding regions
0 lone pairs

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30
Q

Linear bond angle

A

180

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31
Q

Trigonal planar shapes

A

Bonding regions and lone pairs
Trigonal planar:
- 3 bonding regions and 0 lone pairs
Bent/v shape:
- 2 bonding regions and 1 lone pair

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32
Q

Trigonal planar bond angles

A

120

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33
Q

Tetrahedral shapes

A

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

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34
Q

Tetrahedral bond angle

A

109.5

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35
Q

Trigonal bipyramidal shapes

A

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
Q

Trigonal bipyramidal bond angles

A

Trigonal bipyramidal:
90 and 120
See-saw:
180, 120, and 90
T - shape:
180 and 90
Linear:
180

37
Q

Octahedral shapes

A

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
Q

Octahedral bonding angles

A

Octahedral:
90 and 180
Square pyramidal:
90 and 180
Square planar:
90 and 180

39
Q

Dipole

A

Small charge difference across a bond/molecule that results from a difference in electronegativity of the atoms

40
Q

Intermolecular force

A

An attractive force between neighbouring molecules

41
Q

Weak intermolecular forces

A

Temporary dipole-dipole attractions
Permanent dipole-dipole attractions
Hydrogen bonding

42
Q

Temporary dipole-dipole attractions

A

Between non-polar molecules or atoms
AND
Between polar molecules

43
Q

Factors affecting strength of TDD

A

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
Q

Permanent dipole-dipole attraction

A

Between polar molecules

45
Q

Polarity

A
  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
Q

Bond polarity

A

Bond polarity results when there is a difference in electronegativity between two atoms in a covalent bond. Creates dipole due to charge separation

47
Q

Weak intermolecular forces in non-polar molecules

A

Temporary dipole-dipole attraction (TDD)

48
Q

Weak intermolecular forces in polar molecules

A

Temporary dipole-dipole attraction (TDD) and permanent dipole-dipole attraction (PDD)

49
Q

Hydrogen bonding

A

Hydrogen directly bonded to NOF
Also has TDD and PPD

50
Q

Drawing hydrogen bonding

A

https://chemistrytalk.org/what-are-hydrogen-bonds/

51
Q

Weak intermolecular force strengths

A

Strongest -> Weakest
Hydrogen bonding
Permanent dipole-dipoles (PDD)
Temporary dipole-dipoles (TDD)

52
Q

Types of solids

A
  1. Ionic - Ions - Strong Ionic Bonds
  2. Metallic - Atoms - Metallic Bonds
  3. Covalent network - Atoms - Strong Covalent bonds
  4. Molecular - Molecules - Intermolecular Forces
53
Q

Ionic Bond

A

Strong electrostatic force of attraction between cations and anions

54
Q

Metallic bonds

A

Attraction between positive nuclei and sea of delocalised valence electrons

55
Q

Covalent bonds

A

Strong bonds due to shared pair of electrons

56
Q

Intermolecular force

A

Force of attraction between neighbouring molecules

57
Q

Intramolecular bonds

A

Bonds with an molecule
Ionic and convalent

58
Q

Bonds type and relative strengths

A

Ionic 1000
Covalent 1000
Hydrogen bonding 50
PDD 10
TDD 1

59
Q

Atoms in hydrogen bonding

A

Period 2
Small atomic radii
Highly electronegative
Have lone pairs of electrons which can attract the partially charged hydrogen of another molecule

60
Q

Temperature

A

The average kinetic energy of particles in a system

61
Q

Temperature vs Heat

A

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
Q

Melting point

A

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
Q

Boiling point

A

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
Q

Latent heat

A

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
Q

Graph explaination

A

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
Q

Solubility

A

The maximum amount of substance that dissolves per litre of solution

67
Q

Attractions in solubility

A

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
Q

Dissolving and polarity

A

Polar dissolves in polar
Non-polar dissolves in non-polar

69
Q

Enthalpy

A

The energy in a substance due to kinetic energy of particles and potential energy in chemical bonds.

The difference in Hp and Hr

70
Q

Exothermic

A

Releases heat energy
-ve change in enthalpy
Forming bonds
Products have less energy than reactants

71
Q

Endothermic

A

Absorbs heat energy
+ve change in enthalpy
Breaking bonds
Products have more energy than reactants

72
Q

ΔrH

A

The enthalpy change when products form from their constituent reactions under standard conditions

73
Q

ΔfusH

A

The amount of heat energy required to change one mol of a substance from solid to a liquid

74
Q

ΔvapH

A

The amount of heat energy required to change one mole of a substance from liquid to gas

75
Q

ΔsubH

A

The amount of energy required to turn one mol of substance from solid to gas

76
Q

ΔcH

A

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
Q

ΔfH

A

The enthalpy change when one mole of a substance is formed from its constituent elements under standard conditions

Is zero for elements

78
Q

Standard states

A

Gases: O2, Cl2, H2, noble gases
Liquids: Bromine (Br) and Mercury (Hg)
Solids: The rest

79
Q

Hess’s Law

A
80
Q

Calorimetry

A
  2. Dissolving = total mass for q = mcΔT
    3.
81
Q

Calorimetry accuracy

A
82
Q

Entropy

A

The measure of disorder, relation to the degree of dispersal of matter and energy, and the degree of randomness

83
Q

Spontaneity

A

Determined by the entropy change of the system and surroundings
Overall increase in entropy is spontaneous

84
Q

Factors affecting entropy of system

A
  1. Volume
  2. Molar mass
  3. Temperature change
  4. Molecular complexity
  5. Phase change
  6. Number of moles
  7. Solid dissolving in water
85
Q

Factors affecting entropy of surroundings

A

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
Q

Spontaneity ΔH and Δs

A

file:///var/folders/cc/y64ly0jj47lcnw4t1mcm5ll80000gn/T/TemporaryItems/NSIRD_screencaptureui_AVpIAB/Screen%20Shot%202024-08-30%20at%2010.57.31%20AM.png

Chemistry (2 decks)

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