Unit 4.3 Intermolecular Forces & Metallic Bonding

4.3.1 Types of Intermolecular Forces

Intermolecular Forces (IMF)

• Forces of attraction between molecules that hold them together

• 3 main types:

  • Dispersion forces

  • Dipole-dipole

  • Hydrogen bonding

Dispersion Forces

• Electrons in constant motion in an atom

• Electron cloud dispersion can be asymmetrical at any given point in time, creating an instantaneous dipole

• Attraction between partial negative and positive instantaneous dipoles form a dispersion force

• Strength of dispersion forces increases with difference in electronegativity and electron cloud movement

Key Facts about Dispersion Forces

• Present in all atoms and molecules

• Weakest IMF

• Strength depends on the number of electrons

• More electrons = stronger temporary/instantaneous dipole

• Exception: dispersion forces can be stronger than dipole-dipole if the atom is large enough

• Molecules composed of only C and H can only have dispersion forces

Dipole-dipole Attractions

• Attractive force between molecules with permanent dipoles

• Stronger than dispersion forces

• Only for small molecules with the same number of electrons

Hydrogen Bonding

• Strongest IMF

• Special type of dipole-dipole attraction

• Conditions for hydrogen bonding to occur:

  • A species with a very electronegative atom (O, N, or F) that has a lone pair of electrons

  • A hydrogen attached to the O, N, or F

• Hydrogen becomes partially positively charged and can form a bond with the lone pair on another molecule

Hydrogen Bond Donors and Acceptors

• Every hydrogen bond has two components

• A molecule can be both the donor and acceptor, able to hydrogen bond with itself

• Hydrogen bond acceptor only requires an available lone pair, not a hydrogen atom

Hydrogen Bonding in Water and Ammonia

• Water can form a maximum of two hydrogen bonds per molecule

• Ammonia can form a maximum of one hydrogen bond per molecule

• Number of hydrogen bonds possible is restricted by the number of

4.3.2 Deducing Intermolecular Forces

• The structure and chemical formular of the molecules will indicate the types of intermolecular forces present

  • Structure and Symmetry → Is molecule polar or not? (See Section 4.2.4)

  • Chemical Formula → How electronegative are the elements in the molecule?

    • Helps to tell you polar bonds

    • Also tells you if hydrogen bonds are possible when there is N, O, or F

4.3.3 Properties of Covalent Compounds

• Types of intermolecular forces indicate physical properties of molecular covalent compounds (melting/boiling point, solubility, and conductivity)

Melting and Boiling Point

• Changing the state means overcoming the intermolecular forces

• The stronger the forces, the more energy is needed to break the attraction between molecules

• Substances with low melting and boiling points = “volatile”

• As the intermolecular forces increase in strength:

  • The size of the molecule increases

  • The polarity of the molecule increases

Solubility

• “Like dissolves like” = non-polar substances dissolve in non-polar solvents while polar substances dissolve in polar solvents

• However, as the size of a covalent molecule increases in size at a certain point, their solubility can decrease

  • This is because the polar part of the molecule remains the smaller part of the overall structure (In other words, the ratio of polar to non-polar decreases)

  • Ex. alcohols (ethanol is soluble yet hexanol isn’t)

• Giant Covalent substances don’t dissolve in any solvents

  • This is because the energy required to overcome their strong covalent bonds from the lattice structure is too great

Conductivity

• Usually, covalent substances can’t conduct electricity in solid or liquid states since they don’t have any free-moving charged particles

• Only in some cases, polar covalent molecules can ionize and conduct electricity

• Other exceptions are Giant Covalent Structures and they can conduct electricity because they have delocalized electrons (the free-moving charged particles required for conductivity) (See Section 4.2.5)

4.3.4 Metallic Bonding

Metallic Bonding

• Metal atoms tend to pack together in lattice structures

  • This causes their outer electrons to be able to move freely throughout the entire structure = “delocalizes electrons”

• Once their valence electrons are delocalized, the metals gain a positive charge, which repel each other, keeping the entire structure neatly arranged in a lattice

• Metallic Bonding involves strong electrostatic forces of attraction between the metal centers and delocalized electrons

Properties of Metals

• Metals are malleable

  • This is because when a force is applied, the metal layers can slide over each other (the attractive forces between the metal ions and the delocalized electrons act in all directions)

  • So, when the layers slide, the metallic bonds can re-form in a new shape and the lattice is not broken

• Metallic compounds are strong and hard

• Due to the strong attraction between the metal cations suspended in a sea of delocalized electrons

• This also causes metals to have a high melting and boiling point

Conductivity

• Unlike non-metals, metals are able to conduct electricity when in the solid or liquid state

  • Because in both states, they have mobile electrons that can move around and conduct electricity (remember: electric current = flow of electrons)

Strength of Metallic Bonds

Not all metallic bonds have equal strength; there are several factors that affect it:

1. Charge on the Metal Ion

The greater the charge on the metal ion,

→ the greater number of electrons in the sea of delocalized electrons

→ the greater the charge difference between ions and electrons

→ the greater the electrostatic attraction

→ the stronger the metallic bond

2. Radius of the Metal Ion

Metal ions with a smaller ionic radii exert a greater attraction on the sea of delocalized electrons

→ requires more energy to break

→ stronger metallic bond

4.3.5 Trends in Melting Points of Metals

• An increase in the strength of electrostatic attraction is caused by:

  • Increasing the # of delocalized electrons in each metal atom

  • Increasing the positive charges on the metal centers in the lattice structure

  • Decreasing the size of the metal ions

Melting Points of Metals Across a Period

• Ex. Compare the electron configuration of sodium, magnesium, and aluminum and observe the # of valence electrons (do they increase or decrease?)

Na = 1s²2s²2p⁶3s¹

Mg = 1s²2s²2p⁶3s²

Al  = 1s²2s²2p⁶3s²3p¹

• Since aluminum ions are smaller in radius than magnesium or sodium ions

  • So considering that aluminum has the most electrons AND has the smallest radius, it has a stronger metallic bonding → higher melting point

• So as you go across a period, the metallic bonding is stronger and the melting points increase

Melting Points of Metals Down a Group

• As you go down a periodic group, the size of the metal cations increase, thus decreasing the attraction between the negative valence electrons and the positive metallic lattice, so the melting point decreases

4.3.6 Alloys & their Properties

• Alloys = mixtures of metals (the metals are mixed together physically but are not chemically combined)

  • Alloys can also be a mixture of metals and non-metals (ex. with carbon)

• The different metal ion mix is spread evenly throughout the lattice (not clumped together) and are bound together by their delocalized electrons

• Alloys are able to form due to the fact that metallic bonds are non-directional by nature

So why are Alloys made?

• They have distinct and desirable properties since the cations are structured differently in the lattice

Alloy Properties

• Greater strength,

• Harder

  • Since the mixture of atoms in an alloy are different sizes, this distorts the regular arrangement of cations

  • So the layers in a lattice structure have a more difficult time sliding over each other, causing the alloy to be harder than a pure metal

• Higher resistance to corrosion/extreme temperatures

Common Alloys & their Uses