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Chem HL > S2.3 - The metallic model > Flashcards

S2.3 - The metallic model Flashcards

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1
Q

Define metallic bonding

A

The electrostatic attraction between cations and delocalised valence electrons (in a lattice structure)

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2
Q

What are the 6 typical properties of metals

A

• Lustere
• Sonority
• Malleability
• Ductility
• Electrical Conductivity
• Thermal Conductivity

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3
Q

Explain the trend between electrons conductivity and periods

A

As the number of valence electrons increases , the number of delocalised electrons moving throughout the lattice structure increases. This increases the electrical conductivity of the metal.

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4
Q

Why aren’t ionic compounds thermal conductors

A

Do not have free moving electrons

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5
Q

Why are metals malleable

A

• Layers of ions can slide past eachother
• Bonding remains unchanged
• Can be hammered into thin sheets

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6
Q

What type of bonding in ionic , metallic and covalent

A

Ionic and metallic - Non-directional
Covalent - directional

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7
Q

Describe the melting point in a polar covalent molecule

A

Low

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8
Q

What’s the difference between molecule and a lattice

A

lattice - large collection of atoms, ions or molecules in a highly ordered repeated pattern

molecule - discrete group of two or more atoms

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9
Q

What two factors affect the strength of the metallic bond

A

• radius of the metal cations
• ionic charge

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10
Q

The smaller the radius of the metal ion , the stronger the metallic bond.

Explain why

A

There is a shorter distance between the positive nucleus of the cation and the surrounding delocalised electrons , therefore stronger electrostatic attraction

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11
Q

Why do metals further down the group have a larger cationic radius

A

Due to the presence of an additional energy level

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12
Q

Why does lithium have the strongest metallic bond in the group

A

• Lithium metal ions have a smaller radius , the lithium nucleus will be closer to the delocalised electrons , than the other group 1 metals and will result in a greater electrostatic attraction

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13
Q

The higher the ionic charge , the stronger the metallic bond . Why ?

A

greater charge on the metal ion
greater number of delocalised valence electrons

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14
Q

Why does magnesium have a stronger metallic bond than sodium ?

A

• Ionic radius will be similar, same number of occupied energy levels

• Magnesium has a 2+ charge , it will have a greater number of electrons making up the surrounding delocalised electrons, than sodium. The greater charge difference increase the strength of the electrostatic attraction between metal ions and delocalised electrons.

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15
Q

Why does the melting point decrease as you go down a group in the s block

A

Due to the decrease strength of the metallic bond as the radius of the metal ion increase (weaker electrostatic attraction)

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16
Q

What makes transition metals transition metals

A

partially filled d-sub shell
number of different oxidation states

17
Q

Why is copper hard

A

Copper has both 4s and 3d electrons and delocalised valence electrons , there is an increased attraction between these and electrons and the metal ions in the lattice. Increased strength of metallic bond results in greater hardness

18
Q

Why has transition metals have high melting points and electrical conductivity

A

They have delocalised d electrons that strengthen the metallic bond by increasing electron density in the lattice

19
Q

State the meaning of hybridisation

A

mixing of atomic orbitals to form new orbitals

20
Q

How do you know if a species has delocalised electrons

A

• Double bonds
• can STILL display resonance

21
Q

Explain formation of pi bonds

A

Sideways Overlap of p orbitals

22
Q

Explain how metal’s conduct electricity

A

Bonding electrons are delocalised
able to flow throughout atom

23
Q

Explain how metals are malleable

A

Can be hammered into thin sheets
atoms capable of slipping past eachother

24
Q

Why are metals strong and hard

A

due to strong electrostatic forces between metal ions and delocalised electrons

25
Q

Why are metals good thermal conductors

A

The delocalised electrons carry the increased kinetic energy and transfer it rapidly through the metal

26
Q

why do metals have a strong melting and boiling point

A

due to strong electrostatic forces of attraction between the cations and delocalised electrons , requires large amounts of energy to move

27
Q

What are the two factors affecting strength of the metal

A

The charge of the metal ion
The radius of the metal ion

28
Q

Explain how the charge of the metal ion affects its strength

A

The greater the charge the greater the number of protons in the sea of delocalised electrons and the greater the charge difference between the ions and the electrons

a greater charge difference leads to a stronger electrostatic attraction therefore a stronger metallicbond

29
Q

How does the radius of the metal ion affect the strength

A

metals with smaller atomic radii exert a. greater attraction on the sea of delocalised electrons

The greater the attraction means stronger metallic bond requiring more energy to break

30
Q

Why does melting point decrease down a group

A

As you go down , the size of the cation increases
this decreases the attraction between the outer electrons and metallic lattice
reduction in melting point

31
Q

Definition of Transition Metals

A

An element with an incomplete d-sun shell that can form at least one stable cation with an incompleted d-subshell

32
Q

Explain why transition metals have partially good electrical conductivity

A

3d and 4s sub shells are close in energy the transition metals are able to delocalise their d-electrons to form metallic bonds.

33
Q

Why do transition metals have high mp

A

They have greater electron density (ability to delocalise the d-electrons)
electrostatic attraction between the large positive charge of the cations and the sea of delocalised electrons are strengthened.
Results in higher melting point, more energy is required to overcome them

34
Q

Trends in atomic radii

A

generally decreases across each period
Increase down a group

35
Q

Why would there EVER be an increase in ionisation energy

A

electrons being removed from and increasingly positively charged ion

36
Q

Why is Al first ionisation energy lower than magnesium

A

Aluminium has one more electron shell than magnesium.
The additional electron shell shield the outer electrons from the attraction of the nucleus making it easier to remove an electron from aluminium.

Chem HL (13 decks)

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