Bonding Flashcards
Define Metallic Bonding.
The strong electrostatic attraction of positive metal ions surrounded by delocalised electrons.
What are the main factors that affect the strength of metallic bonding?
- size
- charge
- number of delocalised electrons
- smaller ion = stronger metallic
- higher charge = stronger metallic bond
Why is the bonding in Magnesium stronger than in Sodium?
- Mg has a greater charge of 2+
- Mg has twice as many electrons in the sea of delocalised electrons
- Mg ions are smaller, meaning there is a greater charge density
What is the structure formed from metallic bonding?
Giant metallic lattice.
Outline the conductivity of metals.
- good conductors of heat because the delocalised electrons transfer energy through the metal very efficiently
- good conductors of electricity because delocalised electrons can flow
Why are metals malleable and ductile and what does this mean?
Malleable - can be hammered into shape.
Ductile - can be pulled into wires.
Layers of ions can slide past each other.
How are the melting and boiling points of metals determined?
The stronger the metallic bonds, the higher the melting and boiling points.
What are other properties of metals?
Insoluble, strong, shiny
How should you write an answer for comparing the melting points of metals?
- State the structure and describe the bonding - what it occurs between
- Compare the strength of attraction being broken
- Compare energy required to break attraction in each substance
(Model Answer):
In terms of structure and bonding, explain why Calcium has a higher melting point than Potassium.
- Both have giant metallic lattice structure, with electrostatic attractions between positive metal ions and delocalised electrons.
- Calcium has a greater ionic charge (2+) than potassium (1+) and has more delocalised electrons.
- So calcium has stronger metallic bonds and more energy needed to break them
Which metal has a higher melting point, Sodium or Rubidium? Why?
- Both have giant metallic lattices, with attractions between positive ions and delocalised electrons
- Na+ ions are smaller than Rb+ ions
- The attraction between ion and delocalised electrons is weaker in Rb+
- Bonding in Na+ is stronger, more energy needed to break bonds
Define an ionic bond.
The strong electrostatic attraction between oppositely charged ions.
What is the structure formed from ionic bonding?
Giant Ionic Lattice.
What is the typical melting point of ionic lattices?
High.
A lot of energy is needed to break the strong electrostatic attraction between ions.
Describe the electrical conductivity of ionic compounds.
Solid ionic compounds cannot conduct electricity because the ions are fixed and cannot flow.
When ionic compounds are molten/aqueous, they do conduct electricity because the ions are free to move.
Why are ionic compounds brittle and can shatter easily?
Because if the rows of ions slide past each other, like charges will be next to each other - this causes them to repel and shatter.
What are other properties of ionic compounds?
- Crystalline
- Solid at room temp
- Generally good solubility
Describe and explain the electrical conductivity of Sodium Oxide and Sodium in solid and molten states.
Sodium conducts in solid and molten states because it has delocalised electrons which are always free to move + carry a current.
Sodium Oxide conducts in a molten state only as ions are free to move + carry a current. In solid state, ions can’t move so can’t carry current.
Define a covalent bond.
A shared pair of electrons between 2 atoms.
- between non-metals
- atoms share some of their outer shell electrons, giving them both a noble gas configuration
What are the structures which contain covalent bonding?
Macromolecular:
- strong covalent bonds between atoms
- only in carbon/silicon dioxide
Simple Molecular:
- strong covalent bonds between atoms
- weak intermolecular attractions between molecules
What is a coordinate bond?
A shared electron pair which have both come from the same atom
- show it by drawing an arrow from the atom to the atom its being donated to
Why is the bond angle for shapes with lone pairs smaller than shapes without lone pairs?
Lone pairs repel more strongly than bonding pairs.
Explain how the electron pair repulsion theory can be used to deduce the shape of, and the bond angle in, PF3.
- P has 5 electrons in its outer shell
- With 3 electrons from 3 Flourine atoms, there are a total of 8 electrons in it’s outershell.
- So 3 bonding pairs, 1 lone pair
- Electron pairs repel as far as possible
- Lone pairs repel more than bonding pairs
- Shape based on tetrahedral
- Lone pair reduces bond angle from 109.5 to 107 to give trigonal pyramidal shape
Define electronegativity.
The ability of an atom to attract electron pair density from a covalent bond.
