periodicity Flashcards
How are elements arranged in the periodic table
They are arranged in the order of increasing atomic number
What is a period on a periodic table
The horizontal rows in the periodic table
What is a group on a periodic table
The vertical columns
What is meant by periodicity
The repeating trends in chemical and physical properties
What change happens across each period
Elements change from metals to non-metals
How can the electron configuration be written in short?
The noble gas before the element is used to abbreviate
Define the first ionisation energy
The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of 1+ gaseous ions
What are the factors that affect ionisation energy
- Atomic radius
- Nuclear charge
- Electron shielding
Why does first ionisation energy decrease between groups 2 to 3
- Decrease between 2 to 3 because in group 3 the outermost electrons are in p orbitals whereas in group 2 they are in s orbital, so the electrons are easier to be removed
Why does the first ionisation energy decrease between groups 5 to 6
- Decrease between 5 and 6 is due to group 5 electrons and in group 6 outermost electrons are spin paired with some repulsion, therefore the electrons are easier to remove
Explain what happens to the first ionisation energy across a period
- The ionisation energies increase
- Nuclear attraction increases because the number of protons increases which increases the nuclear attraction
- The atomic radius decreases because the electrons are pulled closer to the nucleus
- Similar shielding and distance from nucleus marginally decreases
Explain what happens to the first ionisation energy down a group
- Ionisation energy decreases
- Down a group atomic radius increases
- The number of electrons increases so shielding increases
- Atomic radius increases so the outer electrons are further away from the nucleus which reduces nuclear attraction
Why do successive ionisation energies increase within each shell?
- Because electrons are being removed from an increasingly positive ion and there’s less repulsion amongst remaining electrons
- So the electrons are held more strongly by the nucleus.
Define metallic bonding
Metallic bondingis theelectrostatic forceof attraction between thepositive metal ionsand thedelocalised electrons
Describe metallic bonding
- The electrons in the outermost shell of a metal atom are delocalised
- The metal cations are electrostatically attracted to the delocalised negative electrons
- They form a lattice of closely packed cations in a sea of delocalised electrons.
What affects the strength of a metallic bond
- The number of protons/ Strength of nuclear attraction
- Number of delocalised electrons per atom
- Size of ion
Describe how metallic bonding explains the properties of metals
- They are malleable and ductile due to no bonds holding specific ions together, the metal ions can slide past each other when the structure is pulled
- Good thermal conductor because delocalised electrons can pass kinetic energy to each other
- Good electrical conductors because the delocalised electrons can move and carry a charge
- Insoluble except in liquid metals because of the strength of the metallic bond.
Describe the properties of a giant metallic structure
- High melting and boiling point due to strong electrostatic forces between positive ions and sea of delocalised electrons
- Insoluble in water
- Good conductor of electricity because delocalised electrons can move through the structure, it can conduct when solid.
- Malleable as the positive ions in the lattice are identical so planes of ions can slide over one another
Define giant covalent lattices
Giant covalent lattices are networks of covalently bonded atoms.
Why can carbon form giant covalent lattices?
Carbon atoms can form giant covalent lattices because they can each form four strong covalent bonds
What are allotropes?
Allotropes are different forms of the same element, in the same state
List the allotropes of carbon
- Diamond
- Graphite
- Graphene
Describe the structure of diamond
- In diamond, each carbon atom is covalently bonded to 4 other carbon atoms
- The atoms are arranged in a tetrahedral shape - a crystal lattice structure.
Describe the properties of diamond
- Diamond has a very high melting point
- Diamond is extremely hard
- A good thermal conductor because vibrations travel easily through the stiff lattice
- It can’t conduct electricity because all the outer electrons are held in localised bonds
- It won’t dissolve in any solvent
Describe the structure of graphite
- The carbon atoms are arranged in sheets of flat hexagons covalently bonded with 3 bonds each
- The fourth outer electron of each carbon is delocalised between the sheets of hearings.
Describe and explain the properties of graphite
- Weak forces between the layers in graphite are easily broken, so the sheets can slide over each other
- Can conduct electricity due to the delocalised electrons in graphite that aren’t attached to a carbon atom and are free to move along the sheets and carry an electrical charge
- Lightweight and strong because the layers are far apart compared to the length of covalent bonds, so it is less dense.
- Very high melting point due to strong covalent bonds in the hexagon sheets
- Insoluble in any solvent because the covalent bonds in the sheet are too strong to break
Describe silicon
Silicon forms a crystal lattice structure with similar properties to carbon, each silicon atom is able to form 4 strong, covalent bonds.
Describe graphene
Graphene is a sheet of carbon atoms joined together in hexagons, the sheet is 1 atom thick making it a 2-dimensional compound.
Describe the properties of graphene
- It is an electrical conductor due to the delocalised electrons that are free to move along the sheet
- Strong due to the delocalised electrons strengthening the covalent bonds between atoms
- Transparent and light because it is a single layer
What can graphene be used for?
- High-speed electronics
- Aircraft technology
- Touchscreens on smartphones
Describe simple molecular structures
- They contain only a few atoms
- The covalent bonds between atoms in the molecule are strong, the melting and boiling points depend on induced dipole-dipole forces between their molecules
- they are weak and have low melting and boiling points
