Elements of Life: Periodicity Flashcards
elements are classified as s, p or d block according to what
According to which orbitals the highest energy electrons are in
How do u work out which elements are in which block
- groups 1-2 = s block
- groups 5-8 = p block
- transition metals = d block
What does the period tell you ( electronic configuration )
The last biggest number in the electronic configuration ( the no. of shells )
As you go across within in a block what increases
The number of electrons within that subshell
Eg. 3rd column along in p block will =
p3 ( 3 electrons in subshell p )
Definition of periodicity
Pattern in properties across a row which is repeated in each row
As you go across the period what happens to the atomic radius ( size of atom )
Atomic radius decrease ( atom gets smaller )
Why does atomic radius decrease ( atoms get smaller ) as you go across the period (3)
- More protons ( greater positive charge )
- Outer electrons in same shell
- Stronger pull due to more protons pulling the electrons closer to the nucleus
( = decrease size as electrons pulled )
What happens to atomic radius ( size of atom ) as you go down a group
Atomic radius increase ( atom gets bigger )
Why does atomic radius increase ( atom get bigger ) as you go down the group (2)
- More electrons
2. More electron shells
Definition of 1st ionisation energy
Enthalpy change ( energy required ) to remove 1 electron from each atom in a mole of gaseous atoms
Equation for 1st ionisation energy
X(g) —> X+ (g) + e-
Equation for 2nd ionisation energy
X+ (g) —> X2+ (g) + e-
What is the general trend across the period for ionisation energy
Ionisation energy increases
Why does ionisation energy generally increase across the period (3)
- More protons ( more +be charge )
- Smaller atoms
- Stronger attraction from nucleus to outer shell
If the ionisation energy is high or low what does it mean
High = hard to remove outer electron
Low = easier to remove outer electron
What is the general trend down a group for ionisation energy
Ionisation energy decreases
Why does ionisation generally decrease as you go down the group ( 3 )
- More shells
- Distance between nucleus + outer shell increases
- More shielding ( from larger no. of inner electrons
= attraction between nucleus and outer electrons decreases
What 3 factors effect ionisation energy
- Nuclear charge - more protons = stronger= higher ionisation energy
- Distance from nucleus- attraction decease with distance = larger distance = weak attraction = low ionisation energy
- Shielding- more shells = more shielding = weaker attraction= low ionisation energy
What is the exception of general ionisation energy in terms of group 2-3
??? Group 3 has a lower ionisation than 2 ( eg Al higher than Mg )
Why does group 3 have a lower ionisation energy than group 2
( i.e. a dip in the graph )
??? Shielding effect
Group 2 = s orbital
Group 3 = p orbital
P orbital higher energy level than s orbital = further away from nucleus = more shielding = weaker attraction to nucleus = easier to remove outer electron = lower ionisation energy
What is the exception of general ionisation energy in terms of group 5-6
??? Group 6 has a lower ionisation energy than group 5 ( although would expect to higher )
Eg. S is lower than P
Why does group 6 have a lower ionisation energy than group 5
( i.e. a dip in the graph )
???
Group 5 = final orbital only has 1 electron
Group 6 = final orbital has 2 electrons
Additional repulsion between newly paired up electrons = lower ionisation energy
Definition of electronegativity
Power of an atom to attract the 2 electrons in a covalent bond
Compare H2 and HCl
- which molecule has a higher electronegativity
HCl
Cl = 17 protons H = 1 proton
Cl is 17x more +ve = has greater power to attract electrons in covalent bond than H
( on diagram the electrons will be closer to Cl than H where as in H2 electron in middle because strength is equal )
What is the general trend of electronegativity across the period
Electronegativity increases as you go across the period
Why does electronegativity generally increase as you go across the period (3)
- More protons
- Smaller atoms
- stronger attraction from nucleus to 2 electrons in covalent bond
What kind of melting/ boiling point do metals have eg. Na, Mg and Al
High melting / boiling point
Why do metals have high melting/ boiling points
Strong metallic bonding
- electrostatic attraction between +ve ions and -ve delocalised electrons
What is the general trend for melting / boiling point of metals across the period and why
Increases
- smaller ions
- higher charge (on metal ion)
- more delocalised electrons
What is the general trend for giant covalent melting/ boiling points eg. Si ( Silicon ) and why
Very high melting/ boiling point
- giant covalent structure
- requires larger amount of energy
- to break many strong covalent bonds
What is the general trend of melting / boiling point for simple molecular
Low melting/boiling point
- simple molecular = weak van Der Waals ( intermolecular) forces
What is the general rule for atoms with Van Der Waal forces ( simple molecular ) eg. S8, P4, Cl2
The molecule with more electrons will have strong van Der Waal forces
Eg. With given examples above S8 = most strong =higher melting / boiling point
What is the general trend for melting/ boiling point for monatomic elements and why
Very low melting / boiling point
- very weak Van Der Waal forces
