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