Bonding Flashcards
Metallic Bonding Summary
- left of periodic table
- giant metallic lattice
- strong electrostatic attraction between positive metal ions and sea of delocalised electrons
Ionic Bonding Summary
- metal + non metal
- giant ionic lattice
- strong electrostatic attraction between negative and positive ions
Covalent Bonding Summary
macromolecular:
- diamond, graphite, silicone, SiO2
- covalent bonds between atoms
simple molecular:
- covalent between atoms
- weak intermolecular forces between molecules
Metallic Bonding
Strong electrostatic attraction of positive metal ions surrounded by a sea of delocalised electrons
Comparing the strength of metallic bonds
Bonding in Mg stronger than Na
*Mg has a greater charge of 2+
*Mg has twice as many e- in sea of delocalised e-
*Mg ions smaller, meaning greater charge density
*therefore attraction between Mg2+ ions and the delocalised e- is stronger
Properties of metals
Conductivity- all metals good conductors, delocalised e- help transfer energy and can flow
Malleable/ductile- rows of metal ions can slide past each other
Melting/boiling point- stronger they are, higher the BP/MP
What is a covalent bond
shared pair of e- between two atoms
Covalent structure and bonding
Diamond:
* each carbon has 4 covalent bonds
* high MP
* tetrahedral shape
Graphite:
* conducts electricity
* each C has 3 covalent bonds and a delocalised e-
* layers held together by weak intermolecular forces
* high MP
Simple molecular forces
- intermolecular forces act between molecules
- when a simple molecular substance melts or boils it is intermolecular forces that are broken
- IMF are much weaker than covalent bonds so simple molecular compounds have low MP
What is an ionic bond
Strong electrostatic attraction between oppositely charges ions
Ionic structure and bonding
- metal always positive ion
- non metal always negative ion
high MP/BP
Electrical conductivity
Tend to be brittle and shatter easily
Ionic formulae
Coordinate bond
A shared electron pair which have both come from the same atom
behaves same as covalent bond (length, strength)
Valence shell electron pair repulsion theory
strength:
LP to LP > LP to BP > BP to BP
Strength of repulsion determines bond angle
Double bonds => bonding region (one bond drawn)
Shapes of molecules
What is Electronegativity
The power of an atom to attract the pair of electrons in a covalent bond
F = most electronegative
The factors which determine how electronegative an element are:
- The nuclear charge
- The atomic radius
- The shielding
Across Period 2 the electronegativity increases because:
The number of protons increases
The shielding remains the same
Therefore, the ability to attract electrons in a covalent bond increases.
Polarity
Unsymmetrical, difference in electronegativity so produces covalent bond
central atom surrounded by different elements or LP = polar
Intermolecular forces
Hydrogen bonding:
H-F
H-O
H-N
Permanent dipole-dipole:
Polar = difference in electronegativity
Induced dipole-dipole (Van der waals):
Non polar = no difference in electronegativity
Hydrogen Bonding
When it occurs:
* Strongest intermolecular attraction
* Occurs between H (bonded to N, O, or F) and lone pair on a N, O, or F atom on another molecule
How does it arise?
* Very large difference in electronegativity between oxygen and hydrogen
* Creates a dipole on the O-H bond
* Lone pair on Oxygen atom in one molecule STRONGLY attracts a partially positive Hydrogen atom on a different molecule
Permanent Dipole-Dipole Forces
When it occurs:
* Generally weaker than hydrogen bonding
* Occurs between polar molecules
How does it arise?
* Difference in electronegativity leads to bond polarity
* Dipoles don’t cancel therefore the molecule has an overall permanent dipole
* There is an attraction between ∂+ on one molecule and ∂− on another
Induced dipole-dipole (Van der Waals) Forces
When it occurs:
* Generally the weakest force but can be stronger than both hydrogen bonds and p.d.d. if the molecule is large
* Occurs between all molecules but is the importance force for non-polar molecules
How does it arise?
* Random movement of electrons in one molecule (atom) leads to an…
* Uneven distribution of electron, creating a…
* Temporary dipole in one molecule. This…
* Induces a dipole in a neighbouring molecule.
* Dipoles attract
Ice vs water
ice is less dense than water because the H bonds in ice hold the molecules further apart
