Bonding A1 Flashcards
Module 2
What is the definition of a metallic bond?
The strong electrostatic attraction between cations and delocalised electrons.
How does metal ion charge and ionic radius affect the strength of metallic bonding?
The higher the charge, the stronger the bond.
The larger the ionic radius, the weaker the bond.
Describe the structure of metals.
Cations are fixed in position - maintaining structure and shape of the metal.
Each atom donates its outer shell of electrons to a shared pool of electrons which are delocalised.
Delocalised electrons are mobile so move throughout structure (can carry charge).
What is the definition of an ionic bond?
The strong electrostatic attraction between positive and negative ions. (between metal and non-metal)
How does ion charge and ionic radius affect the strength of ionic bonding?
How does strength of ionic bonding affect the m.p/b.p of an ionic compound?
The higher the charge, the stronger the bond.
The larger the ionic radius, the weaker the bond (because the charge is spread over a large surface area).
The stronger the bond, the higher the mp/b.p of an ionic compound.
What do ionic compounds form?
Giant Ionic Lattices
- each ion is surrounded by oppositely charged ions.
What is the definition of a covalent bond?
The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.
What is a lone pair of electrons?
An electron pair not involved in bonding.
What is the average bond enthalpy?
A measurement of bond strength.
In a dot and cross diagram, does there have to be 8 electrons in the outer shell? (Octet Rule)
No there are exceptions - it can be more or less.
Elements in period 3 such as P, S and Cl can hold up to 18 electrons in their outer shell so they can form more bonds.
What happens in a dative covalent compound?
One element gives both electrons to the bond.
What does the electron repulsion theory state?
Electron pairs repel each other to be as far apart as possible.
Put in order the amount of repulsion in terms of electron pairs from greatest to least.
2 lone pairs (repel most)
1 lone pair, one bonded pair
2 bonded pairs (repel least)
What is electronegativity?
The ability of an atom to attract a bonding pair of electrons.
What happens to electronegativity as you go down the group? Why?
It decreases because:
- more shells, so more shielding
- atomic radius increases (outer shell further from the nucleus)
- so valence electrons less attracted to the nucleus.
What happens to electronegativity as you go across a period? Why?
It increases because:
- Nuclear charge increases
- atomic radius decreases
- so valence electrons more attracted to the nucleus.
What are the 3 types of bonds which come from electronegativity differences?
- Non-polar Covalent (electrons shared equally)
- Polar Covalent (electrons shared unequally)
- Ionic (metal gives an electron to a non-metal)
What is a dipole?
Separation of charges
Why are not all molecules with a dipole polar?
Because the molecule is symmetrical so the dipoles cancel out.
What are intermolecular forces? Are they stronger or weaker than chemical bonds?
Attractions between molecules. Much weaker than chemical bonds.
What are the 3 types of intermolecular forces? Describe them.
London (induced dipole-dipole) forces
- Weak attractions that form between all molecules due to uneven distribution of electrons.
Permanent dipole-dipole forces
- Stronger then London forces
- Exists for polar molecules
- Attractions between partially positive and partially negative ends of molecules.
Hydrogen bonding
-Strongest type of intermolecular force
- A special type of permanent dipole-dipole force when there is a O-H, N-H or F-H bond.
- You need a lone pair of electrons for Hydrogen bonds to form.
Describe how London forces form.
1) Uneven distribution of electrons due to chance causes a molecule/atom to have an instantaneous dipole.
2) This induces a dipole in neighbouring molecules/atoms.
3) The positive and negative dipoles from each molecule/atom attract one another, forming a London force.
How does the size of an atom affect strength of London forces? How does this affect m.p/b.p?
The larger an atom, the more electrons, so stronger induced dipole forces.
The stronger these forces, the higher the m.p/b.p.
What do you need to include in your drawing of hydrogen bonding?
- Labelled hydrogen bond (dashed line)
- Lone pair of electrons
- Show dipoles on atoms involved (Delta plus and minus)
Why is water unusual?
It has an unusually high m.p/b.p for a molecule with such a low molar mass.
Why is solid water (ice) less dense than liquid water?
Liquid water: hydrogen bonds constantly broken and reformed.
Solid water: molecules arranged in hexagonal rings to maximise the number of stable hydrogen bonds - not enough kinetic energy to break bonds.
What is the name of the shape of a molecule with 4 bonding pairs and 0 lone pairs of electrons? What is the angle?
Tetrahedral - 109.5 degrees
What is the name of the shape of a molecule with 3 bonding pairs and 1 lone pairs of electrons? What is the angle?
Trigonal Pyramidal - 107 degrees
What is the name of the shape of a molecule with 2 bonding pairs and 2 lone pairs of electrons? What is the angle?
Non-linear - 104.5 degrees
What is the name of the shape of a molecule with 6 bonding pairs and 0 lone pairs of electrons? What is the angle?
Octahedral - 90 degrees
What is the name of the shape of a molecule with 2 bonding pairs and 0 lone pairs of electrons? What is the angle?
Linear - 180 degrees
What is the name of the shape of a molecule with 3 bonding pairs and 0 lone pairs of electrons? What is the angle?
Trigonal Planar - 120 degrees
Which molecule shapes can be symetrical and therefore non-polar?
Linear, Trignal Planar, Tetrahedral and Octahedral
Which elements form giant metallic lattices?
Li, Be, Na, Mg, Al
Which elements form giant covalent structures?
B, C, Si
What elements form simple molecular structures? What are the formulas?
N2, P4, O2, S8, F2, Cl2, Ne, Ar
What is the difference between Graphite, Graphene and Diamond?
ALL made from C atoms and are giant covalent lattices. Allotropes (differernt forms of same element in the same state).
All have v high m.p. and insoluble due to strong covalent bonds.
GRAPHITE
- lots of layers: weak intermolecular forces (London) between layers, strong covalent bonds within layers - 3D
- C atoms arranged in hexagons - 3 covalent bonds
- delocalised electrons (1 electron per C atom)
- lubricant, lightweight, insoluble, conducts electricity…
GRAPHENE
- one sheet of C atoms in hexagons - 2D
- similar structure to graphite with bonding.
- lightweight, transparent, conducts electricity
DIAMOND
- covalently bonded to 4 other C atoms - tetrahedral.
- v hard, good thermal conductor (vibrations travel easily)
- Cannot conduct electricity
(Silicon is also C covalently bonded to 4 other C atoms)
When can ionic lattices conduct electricity?
When dissolved in liquid as ions are free to move.
Are metallic lattices soluble?
No - because strong metallic bonds.
