Unit 2: VSEPR theory

Question 1

VSEPR theory

Answer 1

Valence shell electron pair repulsion theory

Question 2

Atomic radius

Answer 2

Atomic radius: 1/2 the distance between the nuclei of two identical atoms in an element or compound
- Within a group -> atomic radii increase as you go down a group, as outer electrons occupy higher # (and energy) sublevels, farther away from the nucleu
- Within a period: atomic radii decrease across a period; added electrons are in the same energy level, and ^protons means a stronger pull on the electrons in that energy level, thus pulling them in tighter
- Measured in picometers (1pm = 10^-12m)

Question 3

Shielding (IB key term)

Answer 3

Shielding: when electron shells prevent 2 atoms from bonding (electrons repulse each other; making it harder for them to come together to bond)
- Larger ring / more shells -> more shielding
- Stays relatively constant across a period, as # shells stays the same while ^atomic #

Question 4

Ionic radius

Answer 4

Ionic radius: the distance between the nucleus of an ion to the outermost shell
- As electrons are in orbits, the more electrons there are, the larger the ionic radius
- Largest -> smallest radius: anion, neutral, cation
- Decreases across period (L -> R), as ^#protons
Increases as you go down a group, as atoms ^larger

Question 5

Effective nuclear charge

Answer 5

Effective nuclear charge: as an atom has more protons, there is more “pull” so electrons are closer to the nucleus; increases L -> R

Question 6

Electronegativity

Answer 6

Electronegativity: 𝓬𝓱𝓮𝓶𝓲𝓬𝓪𝓵 𝓹𝓻𝓸𝓹𝓮𝓻𝓽𝔂 𝓽𝓱𝓪𝓽 𝓭𝓮𝓼𝓬𝓻𝓲𝓫𝓮𝓼 𝓽𝓱𝓮 𝓽𝓮𝓷𝓭𝓮𝓷𝓬𝔂 𝓸𝓯 𝓪𝓷 𝓪𝓽𝓸𝓶 𝓸𝓻 𝓪 𝓯𝓾𝓷𝓬𝓽𝓲𝓸𝓷𝓪𝓵 𝓰𝓻𝓸𝓾𝓹 𝓽𝓸 𝓪𝓽𝓽𝓻𝓪𝓬𝓽 𝓮𝓵𝓮𝓬𝓽𝓻𝓸𝓷𝓼 𝓽𝓸𝔀𝓪𝓻𝓭𝓼 𝓲𝓽𝓼𝓮𝓵𝓯 (𝓔𝓷)
- As you go across the periods and up the groups, the electronegativity increases
- Increased atom size and decreased effective nuclear charge
- Small peak in the middle of transition metals
- Fluorine is the most electronegative element
- Has a value of four, therefore all elements calculated relative to fluorine

Question 7

Ionisation energy

Answer 7

Ionisation energy: energy needed for removal of the loosest electron that is being held from an atom in the gaseous state
- The stronger the attraction between valence electrons and the nucleus, the higher the ionisation energy
- Across the period -> increase in ionisation energy
- Atom’s radius decreases as we go to the right
- Ionisation energy decreases as we go down a group because the valence electron(s) are further away from the nucleus; thus, hold on the valence electrons becomes weaker

Question 8

Electron affinity

Answer 8

Electron affinity: the energy released when an atom gains an electron
- X(g) + e⁻ -> X⁻(g) + energy
- Electron affinity decreases as we move down a group because the radius of the atom is larger, making the distance between the electron on the valence shell and the nucleus larger, resulting in a weaker attraction
- ^Electron affinity as we move across a period due to the atomic radius decreasing, making the distance between the nucleus and valence electrons smaller, resulting in a stronger attraction

Question 9

Melting point

Answer 9

Melting point: the energy required to change the state of an ion from a solid to a liquid
- The change of state is caused because the energy is high enough to break the bonds of the atom. As such, atoms with stronger bonds (i.e. the ones in the middle) will not break as easily and have a higher melting point

Question 10

Determining polarity

Answer 10

Type of bond, geometry of molecule
- 0, 0.3 -> non-polar covalent
- 0.3, 1.7 -> polar covalent
- 1.7, 4.0 -> ionic bond
- Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral, square planar

Question 11

Non-polar vs polar molecules

Answer 11

Non-polar: no unshared e-pairs
- Central atom surrounded by identical atoms, all which have same electronegativities
Polar: unshared e-pair
- Central atom has at least 2 different atoms surrounding it (which have different electronegativities)

Question 12

Ionic

Answer 12

Ion: an atom with an electrical charge
- Formed from attraction between cations, anions
- Achieves stable octet (eight valence electrons)

Question 13

Covalent

Answer 13

• Formed when atoms share electrons
• Diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2

Question 14

Metallic

Answer 14

Metallic bond: results from the attraction between positive metal ions and the surrounding sea of mobile electrons
- Metals give up electrons readily, but do not have high ENs necessary to attract electrons
- In a pure chunk of metal, valence electrons are delocalized: they do not have one single atom, but are free to move around
- V,Es are shared equally by all atoms, resulting in an electron sea

Question 15

Dipole-dipole forces

Answer 15

• Strongest intermolecular forces; attraction between polar molecules
• Equal but opposite charges separated by a short distance create a dipole
• Dipole arrow points towards negative pole (^EV atom), + tail points towards positive pole

Question 16

Hydrogen bonding

Answer 16

• Particularly strong type of dipole-dipole attraction; strongest of all intermolecular forces
• Attraction between hydrogen atom bonded to a strong e-neg atom (N, O, F) and unshared pair of electrons on another strongly e-neg atom (N, O, F)
• Attration usually represented by a dashed line

Question 17

Van der Waal’s forces / LDF

Answer 17

Intermolecular attractions resulting from constant motion of electrons and the creation of instaneous dipoles and induced dipoles
- ^Electrons -> ^LDF
- Forces are very weak
- Act between all molecules; only intermolecular forces acting on noble gases, non-polar molecular compounds

Question 18

Crystal type: ionic

Answer 18

Lattice points: ions
Bonds between particles: ionic bonds (600-4000 kJ/mol)
Melting points: high
Electrical and thermal conductivity: low in liquids, high in solids
Ductility and malleability: brittle
Hardness: hard
Solubility: generally soluble in polar solvents
Examples: KCl, MgO

Question 19

Crystal type: covalent network

Answer 19

Lattice points: atoms
Bonds between particles: covalent bonds (300-800 kJ/mol)
Melting points: very high
Electrical and thermal conductivity: low (w/some exceptions)
Ductility and malleability: none
Hardness: very hard
Solubility: insoluble
Examples: diamond quartz (SiO2)

Question 20

Crystal type: metallic

Answer 20

Lattice points: cations
Bonds between particles: metallic bonds (50 -800 kJ/mol)
Melting points: high
Electrical and thermal conductivity: very high
Ductility and malleability: usually ductile, malleable
Hardness: hard
Solubility: insoluble
Examples: Cu, Pb

Question 21

Crystal type: covalent molecular (polar)

Answer 21

Lattice points: polar molecules
Bonds between particles: hydrogen bonding, dipole-dipole (5-50 kJ/mol)
Melting points: low-intermediate
Electrical and thermal conductivity: low
Ductility and malleability: fragile
Hardness: soft
Solubility: generally soluble in polar substances
Examples: NH3, CH3COOH

Question 22

Crystal type: covalent molecular (non-polar)

Answer 22

Lattice points: non-polar molecules
Bonds between particles: dispersion forces (0-5 kJ/mol)
Melting points: vv low
Electrical and thermal conductivity: vv low
Ductility and malleability: vv fragile
Hardness: vv soft
Solubility: soluble in non-polar, slightly polar substances
Examples: CO2, CH4