Bonding

Atoms can exist alone, but more often they combine into compounds that can do many things. Some compounds are enormous, like DNA. Others are needed for everyday life, like water.

Bonding is all about the interactions of electrons. Some atoms give electrons away. Others take them. Some share them. Different situations may call for the electrons to do different things. It's all a matter of chemistry.

There's a lot more to it than simple rules, but these basics are the foundation.

There are three types of bonds to know, and they are based on the metallic properties of the atoms. Remember, metals are to the left of the stairs (on the periodic table). Non-metals are to the right of the stairs. Semi-metals (or metalloids) touch the flat parts of the stairs.

Type of Bond	Types of Elements	Description of Electrons	Other Information
Metallic	Metal and Metal	Sea of mobile electrons	Allows for electricity and magnetism
Ionic	Metal and Non-Metal	Transfer of Electrons	Metals give electrons (become "+" cations) Nonmetals take electrons (become "-" anions) Semi-metals (metalloids) will generally form IONIC bonds.
Covalent	Non-Metal and Non-Metal	Sharing of Electrons	Can be single, double, and triple bonds
None	Anything with a Noble Gas	Noble Gases already have a full outer shell	Noble Gases Do Not React (under normal conditions)

Practice Problems

sodium and chlorine	IONIC (metal and non-metal)	lithium and calcium	METALLIC (metal and metal)
carbon and oxygen	COVALENT (non-metal and non-metal)	magnesium and argon	NO BOND (metal and noble gas)
iron and oxygen	IONIC (metal and non-metal)	molybdenum and chlorine	IONIC (metal and non-metal)
lithium and fluorine	IONIC (metal and non-metal)	nitrogen and carbon	COVALENT (non-metal and non-metal)
copper and zinc	METALLIC (metal and metal)	bromine and bromine	COVALENT (non-metal and non-metal)
neon and helium	NO BOND (noble gas and noble gas)	uranium and oxygen	IONIC (metal and non-metal)

Drawing Ionic Bonds

Start with Lewis dot diagrams of the elements that are bonding.
Move electrons from the metal to the nonmetal.
Write the positive charge to the top right of the metal.
- This is equal to the number of electrons lost.
Place the nonmetal in square brackets.
Outside the brackets in the top right corner, write the negative charge.
- This is equal to the number of electrons it gained.

Example 1: Potassium (K) and chlorine (Cl)

Example 2: Calcium (Ca) and oxygen (O)

Drawing Molecular Bonds

Start with Lewis dot diagrams of the elements that are bonding
For molecular (covalent) bonds, show shared e^- with dashes
- H and H become H-H
- O and O become O=O
- H and O become H-O-H

Bond Polarity

Polarity occurs when a molecule has slightly charged ends, or is in an unbalanced structure
Electronegativity: the willingness or ability of an element to draw electrons toward it
- Fluorine has the greatest EN (4.0)
- Francium has the least EN (0.7)
- Reference Table S shows electronegativities
In a bond, a partial negative charge forms on the element with higher EN
A partial positive charge forms on the element with a lower EN
Larger differences in ENs means higher degree of polarity
When two nonmetals with different electronegativities (EN) bond, they form a polar covalent bond
- Ex: C-H
  - carbon has an EN of 2.6
  - H has an EN of 2.2
  - the carbon side is slightly negative
When two nonmetals with the same electronegativity bond (i.e. two of the same element = diatomic elements) they form a nonpolar covalent bond
- Ex: H-H
- electronegativities are the same

Asymmetrical Molecules

ASYMMETRICAL molecules are POLAR because the pulls and pushes DO NOT cancel out
- Ex: H₂O, NH₃, HCl

Symmetrical Molecules

Completely SYMMETRICAL molecules are NONPOLAR because the pulls and pushes DO cancel out
- Ex: CO₂, CH₄, diatomic elements (H₂, O₂, etc.)

Intermolecular Forces

When a molecule is polar, it leads to intermolecular forces
- these are attractions between molecules
In water, for instance, there is a negative pull toward oxygen and a positive shift on hydrogen
- this means hydrogens from one molecule will be drawn toward oxygens of another molecule
- this is called hydrogen bonding and is rather strong

Oxidation Numbers

an oxidation number is the charge that is formed when an element forms an ionic bond
the charge depends on how many valence electrons the element will either give or take from another element
if an element gives away an electron, it becomes positively charged "+" (cation)
if an element takes an electron, it becomes negatively charged "-" (anion)
YOu can use your Reference Table's Periodic Table to find oxidation numbers.
- Some elements have multiple oxidation states. When in doubt, use the top number.

Example: Calcium has 2 valance electrons. It is easier for calcium to lose 2 electrons than it would be to gain six electrons. Therefore, calcium would get a 2+ charge.

Group #	1	2	3-12	13	14	15	16	17	18
# Valence electrons	1	2	-----	3	4	5	6	7	8
Oxidation #	1⁺	2⁺	-----	3⁺	4⁺ / 4^-	3^-	2^-	1^-	0

Crossing Over

To find out what compound will be made when two elements react, you can either draw the dot diagrams and try to pair up unpaired electrons.
- This can become messy and confusing.
- There is an easier way. All you need is the oxidation number for the elements you are combining.
- Then you cross over the numbers to get the subscripts. Subscripts just tell you how many of that element is in a compound.
- Subscripts are not written with + and - signs.

NOTE: This process is best for ionic compounds, but it will also work for covalent compounds.

Example: hydrogen and oxygen

Hydrogen (group 1) and Oxygen (group 16)	Cross over the numbers (not the +/- signs)	This is what you get	Don't write the number "1" if you have it

Practice problems

What compound forms when the following elements react?

magnesium and chlorine	MgCl₂	potassium and nitrogen	K₃N
rubidium and sulfur	Rb₂S	calcium and bromine	CaBr₂
boron and hydrogen	BH₃	strontium and phosphorous	Sr₃P₂
magnesium and oxygen	Mg₂O₂ (really it's MgO)	beryllium and argon	No reaction; argon is a noble gas

Reverse Crossing Over

Start with a binary compound
Move the subscripts across and up to the opposite element
Metals become + and nonmetals become -
You can use this to determine which group the element is from
- Ex: oxidation 2+ is from group 2
- Ex: oxidation 1- is from group 17

Naming Binary Compounds

To name binary compounds (IUPAC system):
- Name the metal
- Name the nonmetal, but change its ending to “ide”
- For polyatomic ions, don't change the ending
Examples:
- NaCl = sodium and chlorine → sodium chloride
- Ca₃P₂ = calcium and phosphorous → calcium phosphide
- Al₂(SO₄)₃ = aluminum and sulfate → aluminum sulfate (*polyatomic ion)

Polyatomic Ions

Polyatomic ions (P.A.I.)
Ions that are made up of several different elements (“poly” = many, “atomic” = atoms)
Often treated as a single “element”
Reference Table E

NH₄⁺ ammonium	NO₃^- nitrate	ClO₃⁺ chlorate	SO₄^2- sulfate
OH^- hydroxide	CO₃^2- carbonate	C₃H₂O₃^- acetate	PO₄^3- phosphate

Hydrates

Hydrates are molecules that contain water: H₂O
Example: CuSO₄ • 5H₂O
- copper sulfate pentahydrate
C₆H₁₂O₆ (sugar) is a carbohydrate
- There are 6 carbon atoms
- There are 6 “water” molecules

Formula Types

Molecular formula
- Precise ratio of elements in a compound
  - C₄H₈
Empirical formula
- Simplest ratio of elements in a compound
  - C₄H₈ → CH₂
Structural formula
- Shows how bonds are arranged