Thermochemistry Concepts

Question 1

Exothermic

Answer 1

-ΔH (negative)

• Reactants have more energy than products (graph)
• Bond making (g>l>s)
• Heat is released into the system
• Temperature of the surroundings increases

Question 2

Endothermic

Answer 2

+ΔH (positive)

• Products have more energy than reactants (graph)
• Bond breaker (s>l>g)
• Heat is absorbed from of the system (taken out).
• Temperature of the surroundings decreases

Question 3

Lewis structure exceptions to octet rule

Answer 3

• Hydrogen never goes in the middle
• Elements more than atom 14 can have more than 8 electrons
• Hydrogen is happy with just 2e-
• Beryllium is happy with just 4e-
• Boron is happy with just 6e-
• Every other element needs 8e-

Question 4

Why do atoms form covalent bonds

Answer 4

Purpose: Atoms form covalent bonds to achieve a stable electron configuration similar to the nearest noble gas (full outer shell).

Energy: Bond formation releases energy, lowering potential energy and stabilizing the molecule. Breaking bonds requires energy input.

Types of Bonds: Covalent bonds can be single, double, or triple, depending on the number of electron pairs shared. More shared pairs = stronger bond.

Molecular Shape: Covalent bonding influences the molecule’s shape, e.g., methane (CH₄) forms a tetrahedral shape due to four single covalent bonds.

Question 5

Which elements
are exceptions to the usual s, p,
d orbital arrangement?

Answer 5

Chromium (Cr) and Copper (Cu) are exceptions.

Cr: [Ar] 4s¹ 3d⁵

Cu: [Ar] 4s¹ 3d¹⁰

They promote an s electron to d to gain stability from half/full-filled d subshells.

In each case, the 4w prbital contains one electron. This is because the 4s and 3d sublevels lie very close together in energy, and the 3d orbitals being either half full or completely full energetically more stable.

Question 6

describe and explain the
differences in ionic radii
compared to atoms

Answer 6

The size of an atom changes when it forms an ion:

Cations are always SMALLER than atoms (their parent atom). Cation = positive charge, means lost electrons; ratio of protons to electrons is therefore higher, less electrons to protons equals less repulsion; means less shells; smaller size.

Anions are always LARGER than atoms (their parent atom). Anion = negative charge, means gained electrons; ratio of protons to electrons is therefore lower, more electrons to protons equals more repulsion; means more shells; bigger size.

Question 7

describe and explain
whether a covalent atom bond is polar or nonpolar

Answer 7

A bond is polar if there’s a difference in electronegativity, causing unequal electron sharing.
Nonpolar if the atoms have equal electronegativity.

Question 8

list the three types of
intermolecular forces in order of strength

Answer 8

Intermolecular Forces (Strongest to Weakest)

Hydrogen bonding

Permanent dipole-dipole

Instantaneous dipole (Van der Waals)

Question 9

describe and explain
whether a molecule overall is
polar or nonpolar

Answer 9

A molecule is polar if it has polar bonds and an asymmetrical shape, causing net dipole.
Nonpolar if the dipoles cancel due to symmetry.

Question 10

describe hydrogen
bonding and which kind of
molecules have them

Answer 10

Strong attraction between H bonded to N, O, or F and a lone pair on N, O, or F of another molecule.
Occurs in: H₂O, NH₃, HF

Question 11

describe instantaneous
dipole forces (Van der Waals
forces) and state which kind of
molecules have them

Answer 11

Temporary dipoles from uneven electron distribution.

Present in all molecules, but only type in nonpolar substances.

Question 12

describe and explain how
intermolecular forces impact
the boiling point and melting
point of a substance

Answer 12

Stronger forces = higher melting/boiling points.
e.g., hydrogen bonds raise BP/MP; weak Van der Waals = low BP/MP.

Question 13

describe permanent
dipole forces and state which
kind of molecules have them

Answer 13

Attractions between permanent dipoles in polar molecules.
Occurs in: HCl, SO₂, CHCl₃

Question 14

describe whether entropy
is increasing or decreasing for a given reaction

Answer 14

Increase (∆S > 0): solid → liquid → gas, more particles

Decrease (∆S < 0): gas → liquid/solid, fewer particles

Question 15

define the terms
spontaneous and
non-spontaneous

Answer 15

Spontaneous: occurs without added energy (∆G < 0)

Non-spontaneous: requires energy input (∆G > 0)

Question 16

describe and explain
whether a given reaction is
spontaneous or
nonspontaneous

Answer 16

Spontaneous if total entropy increases or if:

Exothermic (∆H < 0) & ∆S > 0
Depends on both ∆H and ∆S (∆G = ∆H – T∆S)

Question 17

Define Hess’ Law

Answer 17

The total enthalpy change for a reaction is the same no matter the path taken, as long as the initial and final conditions are the same.

Question 18

Explaining why metal atoms > cations

Answer 18

Metal atoms are larger because cations lose electrons, reducing electron repulsion and causing the electron cloud to contract.

Question 19

Explaining why nonmetal atoms < anions

Answer 19

Anions gain electrons, increasing electron repulsion and expanding the electron cloud, making them larger than their atoms.

Question 20

Justification of difference in electronegativity for elements in same period

Answer 20

Electronegativity increases across a period due to increased nuclear charge pulling bonding electrons closer.

The number of protons is increasing across the period, so electrostatic attraction between the positive nucleus and valence electrons increases and electrons are drawn closer to the positive nucleus.

Question 21

justification of difference in
electronegativity for elements in same
group

Answer 21

Electronegativity decreases down a group as atomic radius increases and bonding electrons are further from the nucleus.

Valance electrons are further from the positive nucleus, so electrostatic attraction between the positive nucleus and valence electrons decreases.

Question 22

justification of difference in electron shielding for elements in same period

Question 23

justification of difference in electron shielding for elements in same group

Question 24

justification of difference in first
ionisation energy for elements in same
period

Answer 24

It increases across a period due to greater nuclear charge and smaller atomic radius, requiring more energy to remove an electron.

Question 25

justification of difference in first ionisation energy for elements in same group

Answer 25

It decreases down a group because electrons are further from the nucleus and more shielded, so less energy is needed to remove them.

Question 26

justification of difference in atomic radii of elements in same period

Answer 26

Atomic radii decrease across a period due to increased nuclear charge pulling electrons closer. The electron-electron repulsion does increase, but the effect of increasing nuclear charge is greater.

Question 27

justification of difference in atomic radii of elements in same group

Answer 27

Atomic radii increase down a group because more electron shells are added. They have the same electron configuration, so no difference in electron-electron repulsion.

Question 28

relating trend in first ionisation energy to atomic radii (across a period)

Answer 28

As atomic radius decreases across a period, ionisation energy increases due to stronger attraction between nucleus and outer electron.

Question 29

comparing trend in first ionisation energy to electronegativity (down a group)

Answer 29

Both decrease down a group because outer electrons are further from the nucleus and more shielded.

Question 30

attractive forces between particles based on polarity or molar mass

Answer 30

Polar molecules have dipole-dipole forces; non-polar molecules have dispersion forces, which increase with molar mass.

Question 31

attractive forces between particles based on shape

Answer 31

More surface area (longer, flatter shapes) increases dispersion forces due to greater contact between molecules.

Question 32

explaining solubility based on attractive forces solute-solvent

Answer 32

Solubility depends on similar polarity; polar solutes dissolve in polar solvents due to matching intermolecular forces ("like dissolves like").

Question 33

explaining why experimental values of ∆rH may be less negative

Answer 33

May be less negative due to heat loss to surroundings or incomplete combustion/reaction.

Question 34

explaining why Δ vapH° is endothermic

Answer 34

Energy is needed to overcome intermolecular forces and separate particles during vaporisation.

Question 35

explaining why Δ vapH° is > Δ fusH°

Answer 35

Vaporisation requires breaking all intermolecular forces, while fusion only partially disrupts them.

Question 36

explaining how ΔrH° varies depending on state of product

Answer 36

Different states (s, l, g) involve different amounts of energy due to varying intermolecular forces.

Question 37

discussion of spontaneity considering entropy changes of system and surroundings

Answer 37

A reaction is spontaneous if total entropy increases, considering both system and surroundings.

Question 38

discussion of spontaneity considering entropy changes of system and surroundings; ∆rH = +ve

Answer 38

Endothermic reactions require heat input; spontaneous only if system entropy increase outweighs energy input.

Question 39

discussion of spontaneity considering entropy changes of system and surroundings; ∆rH = -ve

Answer 39

Exothermic reactions release heat, increasing surroundings’ entropy, often making reactions spontaneous.

Question 40

Organisation of the periodic table

Answer 40

* Elements with the same number of valence electrons are arranged in VERTICAL groups (e.g. group 13 elements all have atoms with 3 valence electrons). * Elements in each HORIZONTAL row/period have the same number of energy levels (e.g. period 2 elements all have atoms with two energy levels).

Question 41

Atomic sublevel aspects

Answer 41

- Energy levels are given a number called the "principal energy levels" (n). - Lowest energy being 1. The larger the number, the further away from the nucleus those electrons are like to be. e.g. the electrons in principal energy level 1 are closer to the nucleus than those in principal energy level 3. - Electrons fill the lower energy levels first. - All energy levels contain one or more sublevels. Each type of sublevel has a characteristic SHAPE and ENERGY associated with it.

Question 42

Atomic structure

Answer 42

- All matter is made up of tiny particles called atoms (no charge). - Atoms are made up of subatomic particles called protons (positively charged), and neutrons (neutral), and electrons (negatively charged). - Protons and neutrons are found in the atoms nucleus (centre). - 99.9% of the atoms mass is in the nucleus. - The electrons move rapidly around the nucleus in shells (different energy levels) within the atom.

Question 43

Aufbau principle

Answer 43

"Build it up" in german. Means electrons must be placed into the orbitals with the lowest energy first.

Question 44

Pauli's exclusion principle

Answer 44

In an atom or molecule, no more than 2 electrons can occupy any one orbital (boxes), and if 2 electrons are in the same orbital, they must have opposing spins, (shown by arrows facing opposite directions).

Question 45

Hund's rule of maximum multiplicity

Answer 45

Orbitals of the same energy are filled with single electrons (the electrons have the same spin) first, then paired up. In other words, electrons must fill empty orbitals of the same energy before being paired off.

Question 46

Electrons configuration of ions

Answer 46

For NEGATIVE ions, ADD electrons. For POSITIVE ions, REMOVE electrons. E.g. What is the electron structure of O ² ⁻ ? 1. Count no. of electrons in atom: 8 2. Add or remove electrons due to charge: 8 + 2 = 10 3. Fill sub-levels as for uncharged atom: 1s²2s²2p⁶

Question 47

Electron configuration of transition metal ions

Answer 47

When transition metals atoms form ions, the 4s electrons are removed before the 3d electrons. This is because the 3d electrons repel the 4s electrons and push them further away from the nucleus, making them the first electrons to be lost during ionisation. Example: What is the electron configuration of Ni2+ ? 1. Count number of electrons in atom = 28 2. Fill sub-levels, remembering 4s is filled before 3d = 1s²2s²2p⁶3s²3p⁶3d⁸4s² 3. Count number of electrons to be removed (Ni has + charge so electrons are subtracted)= 2 4. Remove electrons, starting with 4s = 1s²2s²2p⁶3s²3p⁶3d⁸

Question 48

Simplifying notation

Answer 48

- We can shorten the electron configurations of large atoms. - When an atom has a full outer shell, it is placed in group 18. - Elements in Gr. 18 can form a reference point for other electron configurations. Example: The atom Ne (neon) is in group 18. It can be used as a point of reference. * Ne has the electron configuration of 1s²2s²2p⁶ * we can use Ne as the starting point and then add more electrons. E.g. For Mg (magnesium): [Ne]3s² This also means, with higher energy atoms (which have longer electrons configurations), you can write out their configurations faster by using a condensed version of another atom, and building upon it.

Question 49

What does size of ion depend on?

Answer 49

1. No. of occupied energy levels. 2. number of protons in the nucleus (vs number of total electrons). 3. Shielding of the outermost valence electrons by the inner energy levels from the positive nucleus. This reduces the attraction between the valence electrons and the nucleus. 4. electron-electron repulsion.