Unit 4.1 Ionic & Covalent Bonding

4.1.1 Forming Ions

• Metals are on the left of the Periodic Table

  • Lose electrons from their valence shell

  • Form positively charged cations

• Non-metals are on the right of the Periodic Table

  • Gain electrons

  • Form negatively charged anions

• Ionic Bonds = bonds formed from the transfer of electrons from a metallic element to a non-metallic element

  • Usually results in both the metal and non-metal having a full outer shell

  • Thus, these atoms have electronic configurations that are the same as a noble gas (elements with full outer shells)

Example:

Na (Sodium)

• It is a metal in group 1, so it has 1 valence electron

• After it loses this electron, it becomes Na⁺ (Sodium Cation)

Cation = A positively charged ion

• A sodium ion has the same electronic configuration as neon: [2, 8]

CL (Chlorine)

• It is a non-metal in group 17, so it has 7 valence electrons

• After it gains an electron from a different atom, it becomes Cl^- (Chloride Anion)

Anion = A negatively charged ion

• A chloride ion has the same electronic configuration as argon: [2, 8, 8]

Electrostatic Attractions = Attractions formed between the oppositely charged ions (cations and anions) to form ionic compounds

• Very strong, so it takes a lot of energy to break

• This is why ionic compounds have high melting points

4.1.2 Ionic Compounds

Ionic Lattices

Ionic Lattice = an evenly distributed crystalline structure formed by ions

• Ions arranged in a regular repeating pattern due to electrostatic forces of attraction between cations and anions

• This causes positive charges to cancel out negative charges, causing the final, overall lattice to be electrically neutral

Properties of Ionic Compounds

• Different types of structure/bonding → different physical properties (ie. melting/boiling points, electrical conductivity, and solubility)

Ionic Bonding & Giant Ionic Lattice Structures

• Ionic compounds are strong

  • Their strong electrostatic forces keep ions held together strongly

• Ionic compounds are brittle, so ionic crystals can split apart

• Ionic compounds have high melting and boiling points

  • This is because of their strong electrostatic forces between the ions that keep them held strongly together

  • As charge density of the ions increase, the melting and boiling points increase

    • This is due to the greater electrostatic attraction of charges

    • ex. Mg²⁺O^2- has a higher melting point than Na⁺Cl^-

• Ionic compounds are soluble in water because they can form ion-dipole bonds

• Ionic compounds only conduct electricity when molten or in solution

  • In those two cases, the ions can freely move around and conduct electricity

  • However, as a solid, the ions are fixed, so they are unable to move around

4.1.3 Formulae & Names of Ionic Compounds

Review:

• Ionic compounds are formed from a metal and a nonmetal bonded together

• Ionic compounds are electrically neutral (positive charges = negative charges)

Charges on Positive Ions

• Positive Ions:

  • All metals

  • Some non-metals

    • ex. NH_4⁺ (Ammonia) and H⁺ (Hydrogen)

Charges of ions depend on their position in the Periodic Table:

• Metals in Groups 1, 2, and 13 = charges of 1+, 2+, and 3+

• Charge on ions of the transition metals can vary which is why Roman numerals are often used to indicate their charge

  • ex. Copper (II) Oxide = copper ion has a charge of 2+

  • ex. Copper (I) Nitrate = copper ion has a charge of 1+

Non-metal Ions

• Non-metals in Groups 15-17 have a negative charge and have the suffix “-ide”

  • ex. Nitride, Chloride, Bromide, Iodide

• Elements in Group 17 = gain 1 electron → have a 1- charge

  • ex. Br^-

• Elements in Group 16 = gain 2 electrons → have a 2- charge

  • ex. O^2-

• Elements in Group 15 = gain 3 electrons → have a 3- charge

  • ex. N^3-

• Additionally, there are more polyatomic or compound negative ions (negative ions made up of more than one type of atom)

7 Common Polyatomic Ions

4.1.4 Covalent Bonds

Covalent Bonds

• Covalent Bonding occurs between 2 non-metals

• Involves the electrostatic attraction between the nuclei of 2 atoms and the electrons of their outer shells

• No electrons transferred, but only shared

• 2 atomic orbitals overlap and a molecular orbital is formed (see left side of image, “bonding”)

• Covalent bonding occurs due to the fact that electrons are more stable when attracted to 2 nuclei rather than only 1

  • Why are they more stable?

    • Because sharing electrons allows each of the atoms to achieve an electron configuration similar to a noble gas (octet rule)

• Usually, each atom provides one of the electrons in the bond

• A covalent bond is represented by a short straight line between 2 atoms

  • ex. H-H

• Note: Think about covalent bonds as electrons being in a constant state of motion, or “charge clouds” rather than an electron pair in a fixed position

Predicting Covalent Bonding

• Use differences in electronegativity to predict whether a bond is either covalent or ionic

Electronegativity & Covalent Bonds

• Electron density in diatomic molecules are shared equally

  • ex. H₂, O₂, Cl₂

Coordinate Bonds

• In simple covalent bonds, the 2 atoms involved share electrons

• However, some molecules have a lone pair of electrons that can be donated to form a bond with an electron-deficient atom (an atom that has an unfilled outer orbital)

  • So both electrons are from the same atom

• This type of bonding is called dative covalent bonding or coordinate bond

• An example of a dative bond is in an ammonium ion

  • The hydrogen ion, H⁺ is electron-deficient and has space for two electrons in its shell

  • The nitrogen atom in ammonia has a lone pair of electrons which it can donate to the hydrogen ion to form a dative covalent bond

Multiple Bonds

• Non-metals are able to share more than one pair of electrons to form different types of covalent bonds

• Sharing electrons in the covalent bond allows each of the 2 atoms to achieve an electron configuration similar to a noble gas

  • This makes each atom more stable

• It is not possible to form a quadruple bond as the repulsion from having 8 electrons in the same region between the two nuclei is too great

Bond Length & Strength

Bond Energy

Bond Energy = The energy required to break one mole of particular covalent bond in the gaseous states (in units of kJ mol^-1)

• The larger the bond energy, the stronger the covalent bond is

Bond Length

Bond length = internuclear distance of two covalently bonded atoms

• It is the distance from the nucleus of one atom to another atom which forms the covalent bond

• The greater the forces of attraction between electrons and nuclei, the more the atoms are pulled closer to each other

• This decreases the bondlength of a molecule and increases the strength of the covalent bond

• Triple bonds are the shortest and strongest covalent bonds due to the large electron density between the nuclei of the two atoms

• This increase the forces of attraction between the electrons and nuclei of the atoms

• As a result of this, the atoms are pulled closer together causing a shorter bond length

• The increased forces of attraction also means that the covalent bond is stronger

• Triple bonds are the shortest covalent bonds so they are the strongest

4.1.5 Bond Polarity

• When two atoms in a covalent bond have the same electronegativity the covalent bond is nonpolar (In other words, the difference in electronegativity = 0)

• When two atoms in a covalent bond have different electronegativitiesthe covalent bond is polar and the electrons will be drawn towards the more electronegativeatom

• As a result of this:

  • The negative charge center and positive charge center do not coincidewith each other

  • This means that the electron distributionis asymmetric

  • The lesselectronegative atom gets a partial charge of δ+ (deltapositive)

  • The moreelectronegative atom gets a partial charge of δ- (deltanegative)

  

• The greater the difference in electronegativity the more polar the bond becomes

Dipole moment

• The dipole moment is a measure of how polar a bond is

• The direction of the dipole moment is shown by the following sign in which the arrow points to the partially negatively charged end of the dipole:

4.1.6 Lewis Structures

• Lewis structures are simplified electron shell diagrams and show pairs of electrons around atoms.

• A pair of electrons can be represented by dots

  The Octet Rule = The tendency of atoms to gain a valence shell with a total of 8 electrons

Steps for drawing Lewis Structures

1. Count the total number of valence

2. Draw the skeletal structure to show how many atoms are linked to each other.

3. Use a pair of crosses or dot/cross to put an electron pair in each bond between the atoms.

4. Add more electron pairs to complete the octets around the atoms ( except H which has 2 electrons)

5. If there are not enough electrons to complete the octets, form double/triple bonds.

6. Check the total number of electrons in the finished structure is equal to the total number of valenceelectrons

Incomplete Octets

• For elements below atomic number 20 theoctet rule states that the atoms try to achieve 8 electrons in their valence shells, so they have the same electron configuration as a noble gas

• However, there are some elements that are exceptions to the octet rule, such a H, Li, Be, B and Al

  • H can achieve a stable arrangement by gaining an electron to become 1s², the same structure as the noble gas helium

  • Li does the same, but losing an electron and going from 1s²2s¹ to 1s² to become a Li⁺ ion

  • Be from group 2, has two valence electrons and forms stable compounds with just four electrons in the valence shell

  • B and Al in group 13 have 3 valence electrons and can form stable compounds with only 6 valence electrons

Unit 4.2 Resonance, Shapes, and Giant Structures

4.2.1 Resonance Structures

• Some atoms or elements have structures that don’t seem to fit with what you would expect their typical Lewis Structure to be

• This can be explained by the delocalization of electrons

• Delocalized electrons = electrons in a molecule, ion, or solid metal that are not permanently associated with one atom or covalent bond

Example:

Nitrate (V) Ion

• A molecule with 1 double bond and 2 single bonds

• It has 3 possible Lewis Structures where the double bond moves around and is with each of the three oxygens

• Since there are different possibilities, these Lewis Structures are also called Resonance Structures

• So you would expect each of these Resonance Structures to have explicitly 1 double bond and 2 single bonds, right?

• That’s not the case, however, because studies of the electron density and bond length show that all 3 bonds are equal in length

• In fact, the electron density is spread evenly between the three oxygen atoms

  • The actual bond length for all of them is somewhere between a single and a double bond

  • The actual structure is something between all the resonance structures and is called a resonance hybrid

Resonance Structures of the Nitrate (V) Ion

Steps to Determine a Lewis Structure:

1. Count the # of Valence Electrons

2. Consider the Charge

   • Add more electrons for negative charges

   • Subtract electrons for positive charges

N + 3O + 1

5 + (3 × 6) +1

= 24 electrons

3. Draw a Skeleton

   • Put single bonds between atoms first

   • Generally, the least electronegative atom goes in the center

   • On the ends, put Hydrogens (because Hydrogen can only have 2 valence electrons and can only bond once) and Halogens

4. Subtract Skeletal Electrons from Valence Electrons

   • Use the remaining electrons to create Lone Pairs

   • OR additional bonds (double/triple) as needed

   • Remember the overall goal is to satisfy the Octet Rule

• 3 structures are possible for Nitrate (V) Ion, with 1 double and 2 single bonds

• The negative charge is distributed throughout the ion and is depicted with the negative sign outside of the resonance brackets

• Electron pairs rapidly oscillate between different positions, never really staying a single or a double bond for long

• Criteria for forming resonance hybrid structures: molecules must have a double bond that is capable of migrating from one part of a molecule to another

• In other words, when there are adjacent atoms with equal electronegativity and lone pairs of electrons that can move to another position in order for the double bonds to be in other positions

  • Ex. Carbonate Ion, Benzene, Ozone, and the Carboxylate Anion

4.2.2 Shapes of Molecules

• VSEPR (Valence Shell Electron Pair Repulsion) Theory = A theory that predicts molecular shape and the angles between bonds based on the concepts:

  1. All electron pairs and all lone pairs arrange themselves as far apart in space as possible

  2. Lone Pairs repel more strongly than bonding pairs

  3. Multiple bonds (double/triple) behave as single bonds

• Domains = The regions of negative cloud charge

• Steric Number (SN) = # of atoms + # of Lone Pairs around the central atom (same concept as Domains)

• Can also be denoted as:

  • A = Central Atom

  • B = Bonded Pair

  • E = Lone Pair

• ex. SN = 2 → AB₂

• ex. SN = 3 → AB₃

• ex. SN = 3 (but one of the domains is a lone pair) → AB₂E₁

Steric Number = 2

• If SN = 2, then the angle between bonds is 180°

• SN = 2 → AB₂

• Molecular Geometry = “Linear”

  • ex. BeCl₂, CO₂, HC≡CH

Steric Number = 3

• If SN = 3, then the angle between the bonds is 120°

• SN =3 → AB₃

• Molecular Shape = “Trigonal Planar”

  • ex. BF₃ and CH₂CH₂ and CH₂O

AB₂E₁

• If one of the electron domains is a lone pair, then the bond angle is slightly less than 120° since lone pairs repulse more, pushing against the other two bonding pairs closer together

  • ex. SO₂

• Molecular Geometry = “Bent Linear”

Steric Number = 4

• If SN = 4, then the angle between bonds is 109.5°

  • E.g. CH₄, NH₄⁺

• SN = 4 → AB₄

• Molecular Geometry = “Tetrahedral”

AB₃E₁

• If one of the electron domains is a lone pair, the bond angle is slightly less than 109.5° due to increased lone pair repulsion

  • ex. NH₃

• Molecular Geometry = “Trigonal Pyramidal”

AB₂E₂

• If 2 electron domains are lone pairs, bond angle also less than 109.5°

  • ex. H₂O

• Molecular Geometry = “Bent”

Summary

4.2.3 Predicting Shapes & Bond Angles

1. Draw Lewis Structure

   • Determine the number of bonding (B) and Lone Pairs (E) around the central atom (A)

2. Apply VSEPR Rules

   • Deduce shape and bond angle

4.2.4 Molecular Polarity

Bond Polarity ≠ Molecular Polarity

• Previously, you learned that bond polarity was determined by the difference in electronegative felt between two bonded atoms

• However, now you can determine if a molecule is polar or not

  • Consider:

    1. The polarity of each bond in the molecule

    2. How the bonds are arranged in the molecule

• Note: Some molecules have polar bonds, yet are overall not molecularly polar since the polar bonds in the molecule are arranged in a way that the individual bond dipole moments cancel each other out

  • ex. CH₃Cl vs. CCl₄

• CH₃Cl

  • Has 4 polar covalent bonds that don’t cancel each other out

  • This means the molecule is polar overall

  • The overall dipole moment is pointing towards the electronegative chlorine atom

• CCl₄

  • Also has 4 polar covalent bonds, BUT the individual bond dipole moments cancel each other out

  • So CCl₄ is a nonpolar molecule

4.2.5 Giant Covalent Structures

Giant Covalent Structures

Covalent Lattices

• Covalent Bonds = Bonds between nonmetals in which electrons are shared between atoms

• Giant Covalent Substances = Sometimes, a substance can't bond like a regular molecule. Instead, the bonds between atoms continue forever, forming a big lattice. There are no separate molecules in this situation, and all the nearby atoms are connected by covalent bonds.

  • ex. C

• Allotrope = Different atomic or molecular arrangements of the same element in the same physical state

• Graphite, diamond, buckminsterfullerene and graphene are allotropes of carbon

Giant Covalent Structures Examples

Diamond

• Diamond is a giant lattice of carbon atoms

• Each carbon is covalently bonded to 4 others in a tetrahedral geometry with a bond angle of 109.5°

• This results in a giant lattice with strong bonds in all directions and causes diamond to be the hardest known substance

  

Graphite

• Each carbon atom is bonded to 3 others in a layered structure of hexagons with a bond angle of 120°

• The spare electron is delocalized and moves around in the space between the layers

• All atoms in the same layer are held together by strong covalent bonds while the different layers are held together by weak intermolecular forces

Buckminsterfullerene

• Contains 60 carbon atoms

  • Each atom is bonded to 3 others by single covalent bonds

• The fourth electron is delocalized so the electrons can migrate throughout the structure

  • This allows for the structure to be a semi-conductor

• Has the same shape as a soccer ball, so it is nicknamed the football molecule

Graphene

• Some substances infinitely covalent bond only in two dimensons, forming only layers

  • ex. Graphene

• Graphene is a single layer of carbon atoms bonded in a repeating hexagonal pattern

  • It is so thin, 1 million times thinner than paper, that Graphene is actually considered 2D

Properties of Giant Covalent Structures

• As always, different structures and bonding types have different effects on the physical properties of substances (ie. melting/boiling points, electrical conductivity, and solubility)

Covalent Bonding & Giant Covalent Lattice Structures

• Giant Covalent Lattices:

  • Very High melting and boiling points

  • Large # of covalent bonds

  • A lot of energy is needed to break the lattice

  • Can be hard or soft

    • Hard (difficult to break their 3D network of strong covalent bonds)

      • Diamond

      • Silicon (IV) Oxide

    • Soft (forces between carbon layers are weak)

      • Graphite

    • (Graphene is strong, flexible, and transparent, making it a very useful material)

  • Insoluble in water (Most)

  • Do NOT conduct electricity (Most)

    • The some that do have delocalized electrons:

      • Graphite

      • Graphene

      • Buskminsterfullerene (semi-conductor)

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