Elements of Life: Periodicity Flashcards

1
Q

elements are classified as s, p or d block according to what

A

According to which orbitals the highest energy electrons are in

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2
Q

How do u work out which elements are in which block

A
  • groups 1-2 = s block
  • groups 5-8 = p block
  • transition metals = d block
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3
Q

What does the period tell you ( electronic configuration )

A

The last biggest number in the electronic configuration ( the no. of shells )

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4
Q

As you go across within in a block what increases

A

The number of electrons within that subshell

Eg. 3rd column along in p block will =
p3 ( 3 electrons in subshell p )

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5
Q

Definition of periodicity

A

Pattern in properties across a row which is repeated in each row

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6
Q

As you go across the period what happens to the atomic radius ( size of atom )

A

Atomic radius decrease ( atom gets smaller )

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7
Q

Why does atomic radius decrease ( atoms get smaller ) as you go across the period (3)

A
  1. More protons ( greater positive charge )
  2. Outer electrons in same shell
  3. Stronger pull due to more protons pulling the electrons closer to the nucleus

( = decrease size as electrons pulled )

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8
Q

What happens to atomic radius ( size of atom ) as you go down a group

A

Atomic radius increase ( atom gets bigger )

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9
Q

Why does atomic radius increase ( atom get bigger ) as you go down the group (2)

A
  1. More electrons

2. More electron shells

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10
Q

Definition of 1st ionisation energy

A

Enthalpy change ( energy required ) to remove 1 electron from each atom in a mole of gaseous atoms

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11
Q

Equation for 1st ionisation energy

A

X(g) —> X+ (g) + e-

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12
Q

Equation for 2nd ionisation energy

A

X+ (g) —> X2+ (g) + e-

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13
Q

What is the general trend across the period for ionisation energy

A

Ionisation energy increases

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14
Q

Why does ionisation energy generally increase across the period (3)

A
  1. More protons ( more +be charge )
  2. Smaller atoms
  3. Stronger attraction from nucleus to outer shell
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15
Q

If the ionisation energy is high or low what does it mean

A

High = hard to remove outer electron

Low = easier to remove outer electron

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16
Q

What is the general trend down a group for ionisation energy

A

Ionisation energy decreases

17
Q

Why does ionisation generally decrease as you go down the group ( 3 )

A
  1. More shells
  2. Distance between nucleus + outer shell increases
  3. More shielding ( from larger no. of inner electrons

= attraction between nucleus and outer electrons decreases

18
Q

What 3 factors effect ionisation energy

A
  1. Nuclear charge - more protons = stronger= higher ionisation energy
  2. Distance from nucleus- attraction decease with distance = larger distance = weak attraction = low ionisation energy
  3. Shielding- more shells = more shielding = weaker attraction= low ionisation energy
19
Q

What is the exception of general ionisation energy in terms of group 2-3

A

??? Group 3 has a lower ionisation than 2 ( eg Al higher than Mg )

20
Q

Why does group 3 have a lower ionisation energy than group 2

( i.e. a dip in the graph )

A

??? 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

21
Q

What is the exception of general ionisation energy in terms of group 5-6

A

??? Group 6 has a lower ionisation energy than group 5 ( although would expect to higher )

Eg. S is lower than P

22
Q

Why does group 6 have a lower ionisation energy than group 5

( i.e. a dip in the graph )

A

???
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

23
Q

Definition of electronegativity

A

Power of an atom to attract the 2 electrons in a covalent bond

24
Q

Compare H2 and HCl

- which molecule has a higher electronegativity

A

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 )

25
What is the general trend of electronegativity across the period
Electronegativity increases as you go across the period
26
Why does electronegativity generally increase as you go across the period (3)
1. More protons 2. Smaller atoms 3. stronger attraction from nucleus to 2 electrons in covalent bond
27
What kind of melting/ boiling point do metals have eg. Na, Mg and Al
High melting / boiling point
28
Why do metals have high melting/ boiling points
Strong metallic bonding | - electrostatic attraction between +ve ions and -ve delocalised electrons
29
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
30
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
31
What is the general trend of melting / boiling point for simple molecular
Low melting/boiling point | - simple molecular = weak van Der Waals ( intermolecular) forces
32
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
33
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