Chemical Formulas

Types of formulasIonic bonding-transfer of electrons

Covalent bonding-sharing of electrons

Determining the types of bonds

• To determine the bond types look at values. (Electronegativity value)

• On chart. (Given)• Bond type electronegativity difference• Covalent 0.0 – 0.5• Polar covalent 0.6-1.7• Ionic bonds 1.8 -4.0

Example

• Determine the type of bond that exist between the following:

• 1. Na and Cl• 2. H and O• 3. C and O

Ionic bonds

• Often formed between metals and nonmetals• Covalent bonds are formed between similar

elements• Binary compounds are those with only two

elements.

Writing Ionic chemical formulas

• 1. Write the symbol and charge of the positive ion.

• 2. Write the negative ion and charge.• 3. DO THEY EQUAL ZERO?• YES –You’re done. Rewrite or erase charges• NO – Criss –cross or find the lowest number

both charges go into and multiply charges to get that number. The number of times you multiply becomes the subscript.

Polyatomics

• Polyatomics –elements that as a group have 1 charge

Ex. SO₄-² = sulfate• If you need more than 1 polyatomic ion, you

must put it in parentheses so subscript multiplies out.

• Ex.• ** Roman numerals give the positive charge.

Examples

Examples

Examples• calcium bromide• Iron (III)sulfide• Potassium fluoride• Lithium carbonate

Formula Mass

• The formula mass of a compound is the total mass of all the atoms.

• The mass of one molecule is expressed in amus.

• The mass of whole moles worth is expressed in grams.

Problems

• Determine the formula mass of

Percent composition

• Determine the percent of each element in the whole formula

• Weight of I kind of element• Total weight of the formula X 100 =• Example : Determine the % composition of

iron(III)chloride.

Practice

• Iron(III)chloride

Terms/additional notes

• Chemical formula is a shorthand way to represent composition of a substance

• Ex. C₁₂H₂₂O₁₁• Empirical formula- simplest ratio of elements in a

compound• Ex. CH₃ is empirical • C₂H₆ is molecular (actual) formula• Subscript – Small number that sits low, behind a

symbol. It tells how many atoms of the element are present.

More info

• In forming bonds, electrons are transferred or shared. This creates ions with resulting charges (positive or negative). These are also called oxidation numbers.

• The following shows the changes as a chemical bond forms.

• Na + Cl → NaCl

Na + Cl →NaCl

• The oxidation number of a free (uncombined) element is zero.

• If an element becomes more positive (sodium) it is “oxidized”. (Oxidation)

• If an element becomes more negative (Chlorine) it is “reduced”. (Reduction)

• These are often referred to as redox reactions

Info cont.

• Transfer of electrons always involve energy changes.

• 1 mole Cl + 1 mole e- → 1 mole Cl- + 83 kcal

• Is this exothermic or endothermic?

Rules for oxidation numbers

• 1. The oxidation number of an atom of a free element is zero.

• 2. The oxidation number of a monatomic ion is its charge.

• 3. The algebraic sum of the oxidations numbers of all atoms in a formula must equal zero.

Rules cont.

4. In compounds, the oxidation number of hydrogen is +1 except in metal hydrides where it is -1. HCl (+1) AgH (-1)

5. In compounds, the oxidation of oxygen is -2, except in peroxides where it is -1 or when it is with very electronegative elements and it is +2.

6. In compounds composed of nonmetals, the less electronegative element is positive and the more electronegative element is negative.

7. The sum of the oxidation numbers of the atoms in the formula of a polyatomic ion is equal to its charge.

Writing names of ionic formulas

• 1. Write the name of the positive ion(metal). If the metal has more than 1 possible charge, you must include the charge as a Roman numeral. Ex. Iron(II)

• Write the name of the negative ion (nonmetal). Change the ending to “ide”.

• A polyatomic keeps its name.• Ex. CaBr is Calcium bromide• CuF is copper(I)fluoride• Na₂SO₄ is sodium sulfate

Molecular compounds

• Covalent bonds are formed when elements share electrons. Often there are more compounds formed between the same two elements. (usually two nonmetals).

• Covalent = molecular• Ex. CO = carbon monoxide• CO₂ = carbon dioxide• *Use prefixes to name the number of atoms

not charge.

Molecular compound prefixes

• Don’t use mono on first word.

• 1. Mono 6. Hexa• 2. di 7. Hepta• 3. tri 8. Octa• 4. Tetra 9. Nona• 5. penta 10. deca

Examples

• CF₄• N₂O• Sulfur hexafluoride• Dinitrogen pentaphosphide

THE MOLE

• A mole is a quantity. • It is equal to 6.02 x 10 ^23 number of

anything. Usually representing very small things like atoms, molecules electrons, formula units etc.

6.02 x 10^23 is also called Avogadro’s number.Where does this number come from?

Where does the mole number come from

• 1 amu= 1/12 th mass of carbon -12.

More mole stuff

• The mass in grams of this number of atoms is called gram atomic mass.

• The gram atomic mass of an atom is numerically the same as the atomic mass of the element.

• So… • The mass of a He atom = 4.0 amus• The mass of a mole of He atoms = 4.0 grams.

Mole stuff cont.

Example: Argon 1 atom argon = 39.9 amu

1 mole argon = 6.02 x 10^23 atoms 1 mole argon = 39.9 grams

1 mole argon gas = 22.4 liters (*gases only at STP)

MOLE CONVERSIONS

• Practice and chart

More conversions

• Convert 0.9 moles to atoms