2. chem bonding

Electrostatic attractions between positive and negative charges

Ability of an atom to attract electrons in a covalent bond to itself

1. Across the period: increase
effective nuclear charge: nuclear charge increases but shielding effect remains constant

2. Down the group: decrease
Effective nuclear charge: constant, as both nuclear charge and shielding effect increases
Atomic radius: increases as number of quantum shells increase

Covalent: similar electronegativity
Ionic: large difference in electronegativity

A lattice where metallic cations and a sea of delocalised electrons are held together by metallic bonding

1. Large amount of energy is needed to overcome strong attraction between metal cations and delocalised electrons

2. Delocalised electrons act as mobile charge carriers that can conduct electrical energy well. Thermal energy causes delocalised electrons to move quickly and transfer thermal energy to other parts of the metal quickly.

3. Layers of ions can slide over one another due to the sea of delocalised electrons preventing repulsion between cations as they move.

4. Depends on how easily particles can be moved from their fixed positions in the solid lattice.
Usually hard though g1 and 2 metals are soft.
Alloys: different sized atoms disrupts the neat arrangement of atoms and they can no longer slide over each other easily, increasing hardness and strength

1. Number of valence electrons contributed per metal atom
More valence electrons = more delocalised electrons = stronger metallic bonding
2. Charge density (ionic charge/ionic radius) of cation. (Small + high charge = stronger ionic bond)

Amount of charge per unit surface area of an ion

Lattice of cations and anions held together by ionic bonding

Strong electrostatic forces of attraction between cations and sea of delocalised electrons

Strong electrostatic forces of attraction between oppositely charged cations and anions

Ionic bonding is non-directional and extends infinitely in all directions, while covalent bonds are directional and only exist between 2 specific atoms

Number of ions that surrounds another ion of the opposite charge is the coordinate number of the central ion

1. Size of ion: a smaller ion usually has a lower coordinate number as less space for oppositely charged ions to be surrounding it
2. Relative charge:
E.g., cation with +2 charge needs twice as many anions as a +1 cations

Heat evolved when 1 mole of a pure solid is formed form its constituent gaseous ions

LE ∝ (q+ x q-)/(r+ + r-)
Higher charge density = larger LE magnitude = stronger ionic bond

1. A lot of energy is required to overcome strong electrostatic forces of attraction between oppositely charged ions

2. Can conduct in molten or aqueous state. Ions can act as mobile charge carriers to conduct electricity
Conc. of aqueous solution increase = electrical conductivity increases
as more ions can carry the charge

3. Ions are held in place with strong ionic bonding. Moving ions out of the position requires a lot of energy to overcome, so they are hard (do not deform easily) and rigid (do not bend easily)

4. If a strong enough force is applied, it will force ions of like charges to move next to each other, causing repulsion of ions and the lattice to shatter.

5. Solubility: usually soluble in water, but some might not be

Formed when two atoms share a pair of electrons which are found in the valence orbitals of the atoms, where they overlap.

1. Bond forned by head-on overlap
2. Bond formed by side-on overlap

S orbitals can only form head-on bonds, while p orbitals can form head-on or side-on overlp

pi bond is a side-on overlap and the degree of overlap is smaller than that of a sigma bond, so sigma bond is stronger

Energy required to convert 1 mole of a covalent bond in gaseous state to form gaseous atoms

1. Bond length: smaller atoms will form shorter and stronger bonds, as overlap between valence orbitals is more effective.

2. Bond order: number of covalent bonds between atoms, as multiple bonds are shorter and stronger than single bonds. As more electrons are being shared, stronger attraction.

Van der Waals radius/diameter

* incomplete octet: elements that are stable w/ fewer than 8 e- in valence shell --> H(2), He (2), Li (2), Be (4), and B (6)

*Expanded Octet: elements in period 3 or greater can hold more than 8 e- in valence shell --> Si, P, S, Cl

*Odd # e-: any element w/ odd # of e- won't be able to distribute e- to give 8 to each atom

Elements period 3 and higher can expand their octets by using the energetically accessible d orbitals (up to 18 valence electrons) but period 2 elements' 3d orbitals are not energetically accessible, and can only have a total of 8 valence electrons

ClF3 and ClF5, as Cl atom can expand its octet to hold more than 8 valence electrons.
ClF4 and ClF6 cannot because it results in an unpaired valence electron
ClF7 cannot as Cl atom is too small to accommodate 7 F atoms around it.

A species with an unpaired electron.

Note: very reactive

A covalent bond where both the electrons comes from the same atom. Donor atom must have 1 lone pair and the receiver needs an empty orbital in its valence shell.

A molecule consisting of two identical molecules joined together

1. Electron pairs around the central atom of a molecule arrange themselves as far apart as possible so as to minimise mutual repulsion

2. The repulsion between lp-lp > lp-bp > bp-bp

Lone pair is only attracted to one nucleus while a bond pair is attracted to both, so the lone pair is shorter than the bond pair and is closer to the central atom, so it repels other electron pairs more than a bond pair.

Linear, 180°

Trigonal planar, 120°

Bent, 118°

Tetrahedral, 109.5°

Trigonal pyramidal, 107°

Bent, 104.5°

Trigonal bipyramidal, 90° and 120°

Seesaw, 120° and 90°

T-shaped, 120° and 89°

Linear, 180°

Octahedral, 90°

Square pyramidal, 90°

the lone pair either is the top or the bottom lp — 4bp — bp

Square planar, 90°
lp — bp — lp

Angle between two bond pairs

Tetrahedral: all bond pairs
Trigonal pyramidal: 1 lone pair
Bent: 2 lone pairs
Lone pairs exert a larger repulsion than bond pairs, causing bond pairs to be pushed closer together and result in slightly smaller bond angles

If central atom more negative (e.g., O in H2O, in comparison to S in H2S), then bond pairs will be closer to O in H2O than to S in H2S, resulting in greater repulsion and larger bond angle.

If terminal atom is more negative (e.g., F in NF3, compared to H in NH3), then bond pairs will be further from N in NF3 than to N in NH3, resulting in less repulsion and a smaller bond angle.

Atomic orbitals mix to form new orbitals called hybrid orbitals which can form sigma bonds, which forms the correct molecular shape for each molecule

After that, describe the bonds:

C=C double bond consists of 1 sp2-sp2 sigma bond formed by head-on overlap of 2 sp2 hybrid orbitals, and 1 p-p pi bond formed by side on overlap of the unhybridised p orbitals.

4 C-H bonds, which are s-sp2 sigma bonds formed by head on overlap of 1 sp2 orbital of C and 1 s orbital of H.

Strength decreases when p character of hybrid orbital increases.
sp > sp2 > sp3

Covalent bond formed by two different atoms, resulting in unequal sharing of electrons due to atoms having different electronegativities. This results in a partial charge separation/dipole

Covalent bond where atoms have the same electronegativity, and electrons are equally shared

When the cation attracts the electrons of the anion, causing the anion's electron cloud to be polarised or distorted.

This is the reason for discrepancies in lattice energy !!!!!

1. Polarising power of the cation
If the cation is small and highly charged, it has a high charge density and high polarising power, and higher tendency of distorting the electron cloud of anions, resulting in more covalent character.

2. Electron cloud of anion
large anions have high polarisability, as their valence electrons are further from and less strongly attracted to its nucleus, so electron cloud is more easily distorted by a cation, resulting in greater covalent character in the ionic bonding.

1) if atoms are the same element
2) if the molecule only has C and H
3) if the individual dipoles balance out (watch for the structure) <- school's method

1) dispersion//instantaneous dipole-induced dipole interactions
2) hydrogen bonding
3) permanent dipole-permanent dipole interactions

electrostatic forces of attraction between partial charges. They are generally weaker than covalent bonds and require less thermal energy to overcome.

they occur due to the fluctuations in electron distribution due to constant movement of electrons, causing an instantaneous dipole in one molecule and inducing a dipole in a nearby molecule, resulting in attraction between those molecules.

1) number of electrons in the molecule:
the larger the electron cloud, the more polarisable the electron cloud

2) surface area of contact between between adjacent molecules
the larger the surface area of contact, the more easily induced dipoles are formed, the stronger the dispersion forces.

note that the surface area of contact depends on the overall shape of molecules

attraction between the oppositely-charged ends of two polar molecules

the intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule

1. electronegativity affects strength. more difference in electronegativity: larger dipole charge, so the hydrogen bonding will be stronger

2. number of H atoms bonded to N, O, F and the number of lone pairs available will affect the extent of hydrogen bonding. e.g., H2O has 2 lone pairs and 2 H atoms, and can form 2 hydrogen bonds each, but NH3 only has 1 lone pair and can only form 1 hydrogen bond each, so H2O has a more extensive hydrogen bonds.

CH4, H2O , HF and NH3 have the same number of electrons and hence similar dispersion forces. But BPs of NH3, H2O and HF are higher as they also have hydrogen bonding, which are stronger than the dispersion forces in CH4 and require more energy to overcome.

(also has higher BP than other hydrides like H2S, HCl etc.)

F is more electronegative than N and hence the electronegativity difference between F and H is larger than between and N and H. Hence, F and H have larger partial charges compared to N and H and will form stronger hydrogen bonds that need more energy to overcome.

H2O has 2 lone pairs and 2 electron deficient H atoms and can form two hydrogen bonds per molecule. however, NH3 only has 1 lone pair and HF only has 1 electron deficient H atom, so they only can form 1 hydrogen bond per molecule. hence, H2O has more extensive hydrogen bonding that requires more energy to overcome.

hydrogen bonding within the molecule. occurs when the electron deficient H atom is in close proximity with the lone pairs of N, O , F atom in the same molecule.

Something with more intramolecular hydrogen bonding will have less extensive intermolecular hydrogen bonds which require less energy to break.

e.g., in water, ethanoic acid (CH3CO2H) forms hydrogen bonds in water, but they form dimers in non-polar solvents like benzene, which is unable to form hydrogen bonds. Hence, acid molecules form dimers instead and their Mr is 120 instead of 60.

In ice, each H2O molecule is hydrogen bonded to 4 other molecules in a roughly tetrahedral structure, causing an open lattice to be formed, with empty spaces between H2O molecules, while the more random arrangement of hydrogen bonds in water takes up less space as the h2O molecules are closer together, As ice occupies a larger volume, it has a lower density.

Hydrogen bonded H2O molecules form an array across the surface of the water, allowing objects you might expect to sink float on it.

Lattice where particles are held together by weak intermolecular forces of attraction.

1) low MP and BP as little thermal energy is required to overcome the weak intermolecular forces of attraction between molecules.

2) soluble in organic solvents. some (like ethanoic acid) are soluble in water as they are able to form hydrogen bonds with water molecules.

3) unable to conduct electricity as there are no mobile charge carriers. However, some substances can react with water to produce aqueous ions, which can conduct electricity (e.g., HCl (g))

3-D structure of atoms, bonded together by strong covalent bonds.

Each sp2 hybridised C atom is covalently bonded to 3 other C atoms, forming a 2 dimensional layer of hexagonal rings. Each C atom has one p orbital that side-on overlaps with each other to form a delocalised pi-bonding system with significant dispersion forces between different layers.

Diamond cannot conduct electricity. As all electrons in each C atom are used to form covalent bonds with other C atoms, diamond has no mobile delocalised electron to conduct electricity.

Graphite can conduct electricity. As each C atom is covalently bonded to only 3 other C atoms, they each have 1 mobile delocalised electron which can move along layers and conduct electricity

Diamond is very hard, strong and non-malleable as it has a rigid structure held together by strong covalent bonds, and a lot of energy is required to move particles from their original position.

Graphite is soft and slippery as there are weak intermolecular forces between layers of graphite, allowing layers to slide over each other when a force is applied.

Both are insoluble in water and organic compounds as attractions between solvent molecules and carbon atoms are not strong enough to overcome the covalent bonds between carbon atoms.

Na, Mg and Al have giant metallic structures and relatively high MPs due to strong metallic bonding between cations and sea of delocalised electrons. MP increases from Na to Al as number of valence electrons increase and increasing charge density

Si has a giant molecular structure and a very high MP as a lot of energy is needed to overcome many strong covalent bonds between Si atoms

P4, S8 and Cl2 have simple molecular structures with weak dispersion forces that take little energy to be overcome

Ar has a simple atomic structure and little energy is needed to overcome weak dispersion forces between atoms

Na, Mg and Al have giant metallic structures with strong metallic bonding between cations and sea of delocalised electrons, which act as mobile charge carriers to conduct electricity. As number of valence electrons and hence number of mobile delocalised electrons increase from Na to Al, electrical conductivity increases.

Si has a giant molecular structure with covalent bonds between atoms, and is a semi-conductor since it is a metalloid.

P4, S8, Cl2 and Ar cannot conduct electricity as they have no mobile ions or mobile charge carriers due to valence electrons held strongly in covalent bonds.

A solute is soluble if the energy released from forming solute-solvent interactions is more than or comparable to the energy absorbed to break solvent-solvent and solute-solute interactions

1. ion-dipole interactions
between a cation and a partial-negatively charged atom of a polar molecule, or between an anion and a partial-positively charged atom of a polar molecule.

2. Hydrogen bonding
between one molecule that has a lone pair on N/O/F and a H atom attached to an N/O/F atom.

3. Permanent dipole-permanent dipole
between 2 polar molecules that are unable to form hydrogen bonds

4. Permanent dipole-induced dipole
between a polar and a non-polar molecule

5.Instantaneous dipole-induced dipole
between two non-polar molecules
also comes with pdpd and pdid

e.g., soluble in water:
the energy released when forming hydrogen bonds between ethanol and water molecules are comparable to the energy absorbed when breaking hydrogen bonds between water and hydrogen bonds between ethanol, so it is soluble

e.g., soluble in water:
NaCl is ionic and can form ion-dipole interactions with water, which are strong an extensive. As the amount of energy released when forming ion-dipole interactions is comparable to energy absorbed to break ionic bonds between Na+ and Cl- and hydrogen bonds in H2O

e.g., insoluble in water
Energy released when forming weaker permanent dipole-induced dipole between water molecules and hexane molecules are less than the energy needed to be absorbed to break hydrogen bonds between water molecules and dispersion forces between hexane molecules.

e.g., insoluble in water
As Ag+ and Cl- have low charge densities, they form weaker ion-dipole interactions with water and less energy is released to form them, and it is insufficient to overcome existing hydrogen bonds between water molecules and ionic bond between Ag+ and Cl- in AgCl.