Bonding and Lewis Symbols

chem 1211, chapter 8

What are bonds?

• forces that hold groups of atoms together, making them function as a single unit

• form when the energy of two atoms together as a unit is lower than the separated atoms

• bonds can be largely, but not always entirely, understood in terms of electrostatic interactions

Describe bond energy

• energy required to break a chemical bond

Describe the types of chemical bonds

• non-polar covalent bonds (very little to no difference in electronegativity)

• polar covalent bonds (moderate difference in electronegativity)

• ionic bonds (large difference in electronegativity)

• the difference in electronegativity can be calculated from the position of elements on the periodic table

Describe ionic compounds

• result when a metal reacts with a non-metal

• electrons are transferred, and the resulting charged particles attract one another through electrostatic interactions

Describe the structures of ionic compounds

• ions are packed together in a “lattice” to maximize the attractions between ions; each ion has multiple immediate neighbors and weaker interactions with more distant neighbors

• not simply pairwise interactions and not molecules

• ionic compounds tend to be brittle and have high melting points

Describe Coulomb’s Law

• quantifies the energy of interaction between a pair of charged particles (e.g. ions)

• q is the charge of each ion, r is the distance between the two nuclei in nm

• this can be used for ions with the same charge and opposite charges; negative sign = attractive force and positive sign = repulsive force

deacribe lattice energy

• indicates how strongly ions are attracted to each other in the solid state; also defined as the energy released when the solid crystal forms from separate ions in the gas phase

• hard to measure directly, but can be calculated; related to Coulomb’s Law

• always exothermic (negative, releasing energy)

describe trends in lattice energy concerning ionic charge

• lattice energies become more exothermic (more negative, favorable) as the magnitude of ionic charge increases; higher charge = stronger attraction and stronger attraction = more exothermic

• ionic charge tends to be the more important of the two factors

Describe ionic sizes in lattice energy

• lattice energies become less exothermic (less negative) with increasing ionic radius; larger ionic radius = weaker attraction and weaker attraction = less exothermic (less negative)

• larger ionic radius means the center of the positive charge is farther away from the negative charge

• ionic charge is more important than ionic size in terms of lattice energy, but ionic size will come into play when two ionic compounds have the same charges

Describe covalent bonding and their Lewis structures

• covalent bonding results when electrons are shared by nuclei

• electron are always shared in pairs; typically shown as a line between bonded atoms

• bonds between nonmetals (including H), a non-metal and a metalloid, or two metalloids tend to be covalent

Describe higher order bonds for a given pair of elements

• single bond (2 electrons) is the longest and weakest

• double bond (4 electrons) = shorter and stronger than a single bond

• triple bond (6 electrons) = shortest and strongest of them all

What are Lewis Symbols?

the single atom to single ion representation of valence electrons

Describe the process of drawing Lewis Symbols?

• write the chemical symbol

• determine the # of valence electrons (look at the last digit of the group #; 1-8)

• fill in the valence electrons around the chemical symbol; singly first, then pair up (Hund’s Rule)

What is true regarding Lewis symbols and Bonding?

Lewis symbols show how many bonds should be formed to achieve an octet from the # of unpaired electrons

Describe Lewis Symbols of Ions and how to draw them

• because of the octet rule, Lewis symbols of ions usually have either eight valence electrons or none

• to draw the lewis symbol of an ion :

1. draw the lewis symbol of the atom

2. take away or add electrons based on the ionic charge

3. put the chemical symbol ad electrons in brackets

4. write the ionic charge as a superscript outside the brackets

How do you draw the lewis structure for an ionic compound?

• write the Lewis structure for the cation(s) and anion(s)

• balance the charges

Describe Lewis structures of covalent compounds

• covalent bonds involve the sharing of electrons to achieve a full valence shell (usually an octet)

• bonding pair: two electrons shared between two atoms

• lone pair: two electrons that ar only on one atom; not involved with bonding

Describe multiple bonds

• we may assign double or triple bonds as necessary to achieve an octet with every atom

• two bonding pairs create a double bond

• three bonding pairs creates a triple bond

What is the formal charge?

• the formal charge is the hypothetical charge on an atom in a molecule if all electrons in each bond were distributed equally

• should be calculated for every atom in the structure

What are the “best” Lewis structures?

• the ones that minimize the magnitude of formal charges

• have negative formal charges on the most electronegative elements, and have positive formal charges on the least electronegative elements

• do NOT have the same sign formal charges on adjacent atoms

How do you calculate formal charge?

• determine the initial amount of valence electrons from the element in questions

• subtract the # of electrons in lane pairs an half of the electrons in each bond

• check your work: the formal charges for every element in a compound should add up to the total charge of the compound (in a natural compound, this should be zero)

What is the terminology regarding Lewis Structures of covalent bonds?

• skeletal structure: all atoms arranged in the order in which they bond to each other

• terminal atom: an atom bonded to only one other atom

• central atom: an atom bonded to two or more atoms

What are the guidelines for Lewis structures of covalent bonds?

• count the total # of valence electrons in the structure

• identify the central atom (often the least electronegative element, or the element that can form the largest # of bonds)

• draw a skeletal structure, connecting atoms by single bonds; subtract electrons used in single bonds from your tally of valence electrons remaining

• complete the octets of terminal atoms first

• if any valence electrons remain to be placed, place them on the central atom

Describe the Lewis Structure conventions

• all valence electrons of all atoms must appear; usually, the electrons are paired

• usually, each atom requires an octet. There are a couple exceptions though

• multiple (double/triple) bonds may be needed to complete octets

• hydrogen atoms are almost always terminal atoms

• lower group # and lower electronegativity are usually in the center

• carbon atoms are almost always central atoms

• generally, structures are compact and symmetrical

Describe Lewis structures with multiple bonds

• if, after placing all valence electrons on peripheral and central atoms as non-bonding pairs, you still have atoms that have not completed their octet, look at using lone pairs to form double or triple bonds