ib chem bonding topic 4/14

why is sea of electrons theory used to describe metallic bonding

he valence electrons are not tightly bound to individual atoms but instead are delocalized and free to move throughout the entire metal lattice. sea gives rise to unique properties associated with metallic bonding like conductivity (electrons move across quickly) and malleability (atoms can slide past each other) these mobile electrons are shared among all the metal ions in the lattice. As a result, they are free to move throughout the structure of the metal.

determine vsepr model quicker (only before period 3)

1. Count the number of valence electrons.

2. Count the number of atoms that bond to the central atom and multiply by 8 to account for full octets on all atoms involved (except for hydrogen, which requires 2).

3. Find the number of lone pairs on the central atom by subtracting the number of valence electrons on bonded atoms (Step 2) from the total number of valence electrons (Step 1).

4. Divide the number of VEs not in bonds (from Step 3) by 2 to find the number of LPs. For example,

• If the number is 0, there are no lone pairs on the central atom.

• If the number is 2, there is one lone pair on the central atom.

• If the number is 4, there are two lone pairs on the central atom.

• If the number is 6, there are three lone pairs on the central atom.

5. Use the Table of VSEPR electron and molecular geometries (Table 1) to determine the VSEPR geometry

bond order

number of apparent bonds between a pair of atoms. the number of bonds in a molecule and divide that by the number of bonding regions 

resonance 

describes delocalized electrons within a molecule, have different ways to write lewis structure.

• true structure is an average of all resonances

C60 fullerene 

NOT a giant covalent structure despite being an allotrope of carbon because it is spherical and consists of discrete molecules rather than a continuous structure of molecules held together

• spherical, made of hexagonal and pentagonal surfaces

• not good conductor because it has weak London forces but some delocalized electons

• can make inclusion complexes to trap stuff

graphene

the single layer component of graphite

• forms thin trigonal planar lattic structure

• excellent conductor. delocalized π electrons are free to move across the entire sheet, contributing to high electrical conductivity. 

• create nanotubes when rolled

Graphite 

layers of hexagonal C sheets - bond angle 120, connected via LDFs. has délocalisation so it conducts electricity 

Allotrope

different geometrical modifications of the same element 

alloys

homogeneous mixture of metals or mix of metals + nonmetal (usually carbon)

• often stronger due to irregular structure. more stable and hard to dent or break because of irregular sized ions cant slide past each other

why are metals highly conductive

electrically: delocalized electrons can freely move

thermally: densely packed nuclei with the electron sea can pass vibrational energy easily like a mosh pit

delocalized electrons

electrons within a molecule that are not confined to a specific bond or atom but are spread out over multiple atoms or regions of the structure. These electrons are not tightly bound to any particular nucleus and are free to move throughout the material.

Delocalization commonly occurs in materials with certain types of bonding, such as metallic bonding - good electrical conductivity

Giant covalent structure

do not form discrete molecules - many covalent atoms to connect the molecules, forming a continuous network. typically hard, insoluble in most solvents, high melting point, doesnt conduct electricity when liquid

coordinate bond

type of bond where 1 atom contributes 2 electrons (a whole lone pair) to the bond - usually to fulfill octet rule. eg. Nh4+, H3O.  sometimes can make dimers eg. AlCl3 x2 = Al2Cl6

metallic bonding

an array of positive ions in a sea of electrons. The metal is held together by the strong forces of attraction between the delocalised electrons and the positive ions.

pi bonds

form from side to side p - orbital overlap

- only in double or triple bonds

sigma bond

covalent bonds formed by overlap of atomic orbitals end-to end/head-to-head. found in single, double and triple bonds

hybridization 

the mixing of atomic orbitals to form new molecular orbitals for bonding. determined by electron geometry 

• 4 e domains = sp3 hybridization

• 3 = sp2

• 2 (linear) = linear

why does like dissolve like

polar - they can form electrostatic interactions with other polars = dissolution

nonpolar = lack of conflicting intermolecular forces to prevent dissolution so why not

hydrogen bonds

strongest of dipdip forces. Occur when H bonds to N, O, F (some of the most electronegative and small elements) 

- also contain dipdip and LDFs - act like perm dipole bond

dipole-induced dipole forces

occur between polar and non polar moleculre 

- when instantaneous/regular dipole nears another molecule, it can induce dipole forces

- This induction occurs because the electric field of the first molecule affects the distribution of electrons in the second molecule. 

dipole-dipole forces

permanent dipole due to unequal share of positive and negative charge (e. neg difference)

- main force between polar molecules

london dispersion forces

weak intermolecular forces cause of temporary dipolarity

• momentary and due to random movement (which causes instantaneous dipoles) of electrons around a molecule (part is pos/neg)

• primary force between diatomics and noble gases but are present in all molecules

• strength increases with electrons - greater probability of instantaneous imbalances in electron distribution, resulting in larger and more frequent fluctuations in electron density. 

• molar mass increase = increase in LDF = boiling point increase cuz more electron density

intermolecular forces vs intramolecular forces

inter = between molecules

intra = within molecules

polarity depends on 2 factors in a molecule: 

1.  presence of polar bonds in the molecule

2. geometry of the molecule

* polar molecules have a net dipole moment (towards a certain atom or area)

polarity

the unequal partial charge distribution between atoms within a compound. 

why do lone pairs cause angles between other bonds to decrease?

occupy a larger region of space around the nucleus of an atom. Bonding pairs are shared between atoms and are often localized between the nuclei = smaller region of electron density. The volume of  the repulsion of the lone electrons and the lack of ownership by another atom makes it stronger than the force of those involved in bonding

bonding domain

only those electron pairs involved in bonding

electron domain

all electron pairs involved in bonding or present as lone pairs

covalent bonds

covalent bonds form through the collision of atoms and electron clouds overlap enough to cause attractive forces to exceed repulsive ones - hence electrons are shared. 

* all valence e's can move around within the whole compound

Ionic bonds (ib definition)