Chapter 7 Flashcards
What do groups in the periodic table show? How does this relate to the elements?
- Groups in the periodic table have the same number of electrons in their outer shell, and similar chemical properties
What is periodicity? Required.
- Repeating trends in the physical and chemical properties of elements across a period
What are 4 properties that have periodicity?
- Electronic configuration
- Ionisation energy
- Structure
- Melting points
What trend do periods have in terms of electronic configuration?
- Each period starts with an electron in the new highest energy level
What is first ionisation energy? Required.
- The energy required for the removal of 1 mol of electrons from 1 mol of gaseous atoms of an element
Write an equation for the first ionisation energy of chlorine.
- Cl (g) -> Cl+(g) +e- (it is not shown as diatomic as first ionisation energy refers to 1 mol of atoms)
What is second ionisation energy?
- The amount of energy needed to turn 1 mol of gaseous 1+ ions into 1 mol of gaseous 2+ ions
What 3 factors affect ionisation energy?
- Atomic radius
- Nuclear charge
- Electron shielding
How does atomic radius affect ionisation energy?
- The greater the distance between the nucleus and the outer electrons, the less the nuclear attraction
How does nuclear charge affect ionisation energy?
- The more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and the outer electrons
How does electron shielding affect ionisation energy?
- Electrons are all negatively-charged and therefore repel each other, meaning that inner-shell electrons repel outer-shell electrons
- This reduces the attraction between the nucleus and the outer electrons
How do successive ionisation energies compare with each other?
- They increase
- There is a big increase in the ionisation energy between different shells of an atom
Why do successive ionisation energies increase?
- The atomic radius decreases
- The nuclear attraction between the remaining electrons and nucleus increases (the same number of protons attract fewer electrons)
Describe the trends in first ionisation energy in the periodic table.
- They decrease down a group (down across the different periods)
- They generally increase across a period
Why does first ionisation energy decrease down a group?
- The atomic radius and shielding increases
- Nuclear attraction on outer electrons decreases
Why does first ionisation energy generally increase across a period?
- The nuclear charge increases, which causes the atomic radius to decrease
- There is similar shielding
- The nuclear attraction therefore increases
In what 2 places does the first ionisation energy not increase across a period?
- Between groups 2 and 3
- Between groups 5 and 6
Why does the first ionisation energy decrease between groups 2 and 3?
- In group 3, the p sub-shell starts to be filled
- As it has more energy than the s sub-shell, it is easier to remove the electron in the p sub-shell
- Group 3 therefore has a lower first ionisation energy than group 2
Why does the first ionisation energy decrease between groups 5 and 6?
- In group 6, the p sub-shell electrons start to pair up
- The paired electrons in one orbital repeal each other, making it easier for one of them to be removed
- Therefore, group 6 has a lower first ionisation energy than group 5
What does the graph of first ionisation energy against atomic number look like? Give 5 details.
- The highest points are the noble gases
- The lowest points are the group 1 metals
- First ionisation energy generally increases across each period
- There are dips between groups 2 and 3 and 5 and 6
- Each period has a lower first ionisation energy than the previous period (the entire line is shifted down)
What is metallic bonding? Required.
- The strong electrostatic attraction between cations and delocalised electrons
What structure do metals have?
- They all have a giant metallic lattice structure
What are the 2 main properties of metals?
- High electrical conductivity
- High melting and boiling points
When can metals conduct electricity, and how?
- When in the solid or liquid states
- This is because the delocalised electrons are free to move through the structure and carry charge
Why do most metals have high melting and boiling points?
- Large amounts of energy are needed to overcome the strong electrostatic attraction between the cations and electrons
Are metals soluble? Why?
- No
- They react with solvents instead, such as group 1 metals with water
What are giant covalent lattices? Required.
- Networks of atoms bonded by strong covalent bonds
List 2 elements that form giant covalent lattices.
- Carbon
- Silicon
What are the melting and boiling points of giant covalent lattices like, and why?
- They are high because a lot of energy is needed to break the strong covalent bonds
Are giant covalent lattices soluble? Why?
- No
- The covalent bonds are too strong to be broken due to interactions with the solvent
Can giant covalent lattices conduct electricity? Why?
- Most aren’t as they don’t have any electrons that are free to conduct electricity
- Graphite and graphene are exceptions to this as they only use 3 of their 4 electrons in their outer shell to form covalent bonds
What is the difference between graphite and graphene?
- Graphene is a single layer of graphite (in graphite the separate layers are bonded together by London forces)
How do melting points vary across periods 2/3?
- The melting points increase from periods 1 to 4
- There is a sharp decrease between groups 4 and 5
- From group 5 to 0, the melting points are low
Why do melting points decrease between groups 4 and 5?
- The elements change from giant molecular structures to simple molecular substances
Why do melting points increase between groups 1 and 3?
- The larger the charge of the cation (and therefore also the relative amount of delocalised electrons) the stronger the metallic bonds
