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Chemistry > 3.1.1: Periodicity > Flashcards

3.1.1: Periodicity Flashcards

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1
Q

Define the term first ionisation energy

A

the energy required/ energy change to remove an electron from each atom in one mole of gaseous atoms to form one mole of gaseous ions

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2
Q

Write an equation for the second ionisation energy of oxygen

A

O+(g) → O2+(g) + e-

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3
Q

Suggest why the second ionisation energy of oxygen has a greater value than the first

A

the O+ ion, is smaller than the O atom
OR the electron repulsion/shielding is smaller OR the same number of protons/nuclear charge attracting a fewer number of electrons

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4
Q

How do ionisation energies mark a new shell ?

A

Large difference between ionisation energies mean a new shell closer to the nucleus

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5
Q

Describe and explain trends across period 2 Li –> N

A
  • Li —>N ionisation energy increases as atomic radius decreases, nuclear charge increases, nuclear attraction increases and electrons experience same shielding
  • Be —> B ionisation energy decreases as electron is removed from p orbital which has a higher energy
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6
Q

Why does nuclear attraction increase

A

Because shell is drawn in closer to nucleus due to increased nuclear charge/ electrons experience greater attraction

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7
Q

What does a sharp rise in successive ionisation energy between 3rd and
4th IE indicate ?

A

The element is Al as it marks a change to a new shell / there are 3
electrons in the outer shell

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8
Q

Explain why the first ionisation energy of B is less than that of Be

A

-In B, electron being removed is at a higher energy
-In Be, electron being removed is at a lower energy
-An s electron is lost in Be AND a p electron is lost in B

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9
Q

Why is the second ionisation energy of Ca greater than the first ?

A

-same number of protons or same nuclear charge attracting
-less electrons and ion is smaller
- Nuclear attraction increases

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10
Q

How is the periodic table arranged ?

A

By increasing atomic number

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11
Q

Why is the periodic table not arranged by atomic mass ?

A

Atomic mass changes due to isotopes

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12
Q

Give 2 physical properties of metals

A

Conduct electricity and are ductile

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13
Q

Give 2 chemical properties of metals

A
  • Form ionic compounds with non metals
  • Form positive ions by electron loss
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14
Q

Give 2 physical properties of non metals

A

Brittle and poor electrical conductor

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15
Q

Give 2 chemical properties of non metals

A
  • Form ionic compounds with metals
  • Form negative ions by electron loss
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16
Q

Give the history of the development of the periodic table

A
  • Ordered by atomic mass
  • TRIADS - grouped into 3s by characteristics
  • NEWLANDS OCTAVES - grouped by mass with every 8th element having similar properties however transition metals didn’t fit patters
  • MENDELEEV - ordered by atomic mass, left gaps for undiscovered elements and predicted properties
  • ORDERED BY PROTON NUMBER - periods and groups
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17
Q

Is ionisation energy endothermic or exothermic ?

A

Requires energy so is endothermic (positive value)

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18
Q

Give the 3 factors affecting ionisation energy

A

Atomic radius, nuclear charge, shielding

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19
Q

Where are the decreases in ionisation energy in period 3 ? Why?

A
  • Mg and Al as the outer electron in Al is in a higher energy subshell (p)
  • P and S as electron repulsion in S is higher as removing from S when 2 electrons are in an orbital requires less energy than 1 electron being in an orbital
20
Q

Where are the decreases in ionisation energy across period 2 ?

A
  • Be and B
  • N and O
21
Q

Give the reasons for the differences in melting points across periods 2 and 3

A

First 3 elements have metallic bonding with increased positive charge and no. of delocalised electrons forming a stronger metallic bond. Giant covalent structure. Simple molecular structure. Larger simple molecular structure. Smaller simple molecular structure. Exists as individual atom.

22
Q

What are the 3 giant covalent structures ?

A

Graphite, diamond and graphite

23
Q

Describe features of graphite

A
  • Layers can slide as weak forces between layers
  • Each carbon bonded 3 times with 4th delocalised electron
  • High electrical conductivity as delocalised electrons between layers
  • Lots of strong covalent bonds causing high mp
  • Insoluble as covalent bonds are too strong to break
  • Low density as layers are far apart compared to length of covalent bonds
24
Q

Describe features of diamond

A

-Each carbon bonded 4 times in tetrahedral shape
- Good thermal conductor as highly packed
- High melting point meaning strong covalent bonds
- No delocalised electrons meaning it doesn’t conduct electricity
- Insoluble as covalent bonds are too strong to break

25
Q

Describe features of graphite

A
  • 1 atom thick layer of graphite with hexagonal carbon rings
  • Good electrical conductor as free moving electrons
  • Lightweight, transparent due to strong covalent bonds (used in smart phone screens)
26
Q

What is silicon the same as ?

A

Diamond

27
Q

What is the structure of compounds with metallic bonding ?

A

Giant metallic lattice structure

28
Q

What are features of metallic compounds ?

A
  • Good thermal conductor as delocalised electrons
  • Good electrical conductor as mobile delocalised electrons
  • High melting point as strong electrostatic attraction between bonds
  • Insoluble as strong metallic bonds
29
Q

Describe how bromine bonds

A

Simple molecular structure with london forces between molecules

30
Q

Describe bonding in magnesium bromide

A

Ionic bonding forming a giant ionic lattice between oppositely charged ions

31
Q

Describe bonding in magnesium

A

Metallic bonding forming a giant metallic lattice with delocalised electrons

32
Q

Why can Br2 not conduct electricity ?

A

Electrons are not mobile

33
Q

What is melting point linked to ?

A

Strength of bonds/intermolecular forces

34
Q

What type of structure is Si02 and what bonds does it contain ?

A

Giant covalent lattice with covalent bonds

35
Q

When do you mention that increased shielding outweighs increased nuclear charge ?

A

When explaining why ionisation energies decrease DOWN GROUPS relating to reactivity

36
Q

Explain the difference in melting points between magnesium and sodium

A
  • Mg has more outer electrons
  • Magnesium ions have a greater (positive) charge (density)
  • Magnesium has a greater attraction between ions and delocalised electrons
36
Q

Explain differences in melting point between P4 and Cl2

A
  • P4 has more electrons
  • P4 has stronger london forces
  • It requires more energy to break london forces in P4 than Cl2
37
Q

What bonds and forces occur between Silicon ?

A

Covalent between atoms

38
Q

What do london forces occur between ?

A

Atoms

39
Q

How does attraction between nuclei and outermost electrons change across a period?

A
  • Attraction increases across a period and nuclear charge increases (as number of protons increases)
  • Same number of shells (so outer electrons experience similar shielding)\
  • Atomic radius decreases
40
Q

Explain the trend in atomic radius across a period

A
  • Atomic radius decreases and nuclear charge increases
  • Same shielding
  • Greater nuclear attraction ON OUTER ELECTRONS
41
Q

Why do successive ionisation energies increase with ionisation number ?

A

radius decreases AND attraction between (the remaining) electrons and nucleus increases

42
Q

What is the removal of one electron from each atom in 1 mole of gaseous atoms called ?

A

First ionisation energy

43
Q

Explain trends in reactivity down group 7

A
  • Reactivity decreases
  • Atomic radius increases
  • More shielding
  • Nuclear attraction decreases
  • Increased shielding outweighs increased nuclear charge
  • Ionisation energy decreases
44
Q

Why is it unnecessary to refer to carbon as diamond or graphite when referring to ionisation energy ?

A

Diamond and graphite form gaseous atoms of carbon when ionised

45
Q

What are the particles between which the bond is acting in lithium ?

A

Li+ ions and delocalised electrons

46
Q

Define periodicity

A

The repeating patterns in the periodic table

Chemistry (33 decks)

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