1. Chemical Formulas and equations

2. Chemical formulas
Formulas use chemical symbols and numbers for qualitative and quantitative information
Symbol = qualitative (element)
Example, CO
Subscript = quantitative (amount)
Example, CO2

3. Types of formulas
There are two types of formulas, empirical and molecular.
The empirical formula represents the simplest integer ratio in which atoms combine to form a compound.
Ionic compounds do not form molecules, so ionic formulas indicate a ratio of ions.
Example, NaCl (1 sodium and 1 chlorine ion)
Example, MgCl2 (1 magnesium and 2 chlorine ions)
Formulas of ionic substance are empirical

4. Types of formulas
Molecular formulas may be a multiple of the empirical formula.
Covalently bonded substances form molecules.
A molecular formulas represents the actual ratio of atoms in a molecule.
Example, glucose is C6H12O6, but that formula can be reduced to CH2O

5. Atoms, compounds, and ions
Atoms and compound are electrically neutral.
In an atom protons = electrons
In a compound the charges balance out.
Ions are not neutral and may be either positive or negative (cations and anions)
An ionic charge is indicated by a superscript.
A sodium ion with a charge of 1+ is written as Na+ (this is known as the oxidation state)
Some atoms have more than one oxidation state

6. Polyatomic ions
A polyatomic ion is a group of atoms covalently bonded together possessing a charge.
Parentheses are used to enclose polyatomic ions when there is more than one of the ions in a compound.
A subscript after the parentheses tells how many polyatomic ions are present.

7. coefficients
A coefficient written in front of a formula tells how many units of the formula are present and it applies to the entire formula
Example, 2H2O
The elements present are: hydrogen and oxygen
In water there are two hydrogens for every one oxygen
In 2H2O there are actually 4 hydrogen atoms and 2 oxygen atoms
How many of each atom in: 4 Ca(NO3)2

8. What is a Hydrate
A hydrate is formed when water gets trapped with in the crystal lattice structure of an ionic compound.
These crystals have a definite number of water molecules for each unit of the compound
An anhydrous (opposite of hydrated) compound can be obtained by heating the crystals to drive off water.
In a chemical reactions, the water in the hydrate does not react, but it adds mass to the compound.

9. Writing formulas and naming compounds

10. Equalizing charges
Compound are considered neutral by having an equal number of positive and negative charges.
The criss-cross method is used to help figure out the correct formula when equalizing charges.
Remember to transfer only the number and not the sign and do not write the number 1

11. Naming compounds
Compounds are named based on the types of elements that form them.
Ionic compounds are named by one method.
Covalent compounds that contain only nonmetals are named differently.

12. Binary ionic compounds
The names of binary ionic compounds come directly from the elements in the compound.
The positive particle (the metal) is placed first.
The negatively charged ion will end the formula.
The metal’s name remains the same and the nonmetal gets an –ide ending.
Example, NaCl is sodium chloride
Example, CaBr2 is calcium bromide

13. Ionic compounds with polyatomic ions
Naming these kinds of compounds is simple, just simple use the metal’s name and the polyatomic ions name.
Nothing needs to modified.
Example: NaOH is sodium hydroxide

14. Binary Covalent Compounds
A binary compound that contains two nonmetals is arranged by electronegativity values.
The element with the lower electronegative value is written first.
The name of the compound will end in –ide.
These elements can often form more than one compound, so a prefix is used to tell the reader how many atoms of each element are present.
Example, CO is carbon monoxide
Example, CO2 is carbon dioxide
If there is only one atom of the first element in the compound a prefix is not needed

15. The Stock System
Some metals (transition metals) have more than one oxidation state.
The stock system solves this problem by simply stating the oxidation number by using Roman numerals after the name of the metal.
Example, Iron (II) chloride tells the reader that the iron has an oxidation number of +2 and the formula would be FeCl2
What would Iron (III) chloride look like?

16. Chemical Reactions Physical Change
A change in the appearance of the starting material
Phase changes are physical changes
Ex, melting, freezing, boiling
Chemical Change
A change in which the products are a different material than the reactants
Chemical reactions are chemical changes
Ex, burning, rusting

17. Chemical Equations
A chemical equation is used to show what takes place during a chemical reaction.
The starting substance is called the reactant (located on the left side of the arrow)
The substance produced is called the product (located on the right side of the arrow)
H2 (g) + O2 (g)  2H2O (g)

18. Endothermic and Exothermic
Chemical and physical changes involve the loss or gain of energy.
Endothermic – heat is absorbed during the reaction
Exothermic – heat is released during the reaction

19. Type of reaction
Surrounding temperature
Potential Energy of the Reactants
Potential Energy of the Products
Value of Delta H
Endothermic
Decreases
Less
More
Positive
Exothermic
Increases
Negative

20. Endothermic
Process that requires energy in order for a chemical reaction to occur.
Physical change – ice melting
Chemical change – food cooking
The reactants absorb energy as they become products.
The products have more potential energy than the reactants.

21. Exothermic
Process that release energy when a chemical reaction occurs.
Physical change – water freezing
Chemical change – combustion
The reactants release energy as they become products.
The products have less potential energy than the reactants.

22. Balancing chemical reactions
The Law of Conservation of Mass: mass of the products = the mass of the reactants
There also has to be conservation of atoms
Coefficients in equations tells us how many atoms or molecules are present (Moles)

23. Types of reactions Synthesis reactions Decomposition reactions
Single replacement reactions
Double replacement reactions

24. Synthesis reactions
When two or more reactants combine to form a single product
A + B  AB
A and B represent either elements or compounds and AB represents a compound

25. Decomposition reaction
The reverse of a synthesis reaction in that a single compound is broken into tow or more simpler substances
AB  A + B
A and B represent either elements or compounds and AB represents a compound

26. Single replacement reactions
A type of reaction where one element replaces another element in a compound.
This type of reaction always involves an element and a compound.
A + BC  B + AC
One can predict whether a reaction will happen or not using the Activity Series
A metal listed on the table will react with the compound of a metal below it.
A nonmetal will replace a less active nonmetal in a compound

27. Double replacement reaction
These reactions generally involve two soluble ionic compounds that react in solution to produce a precipitate, a gas, or a molecular compound
AB + CD  AD + CB
There are three ways to determine if a double replacement reaction will occur

28. Double replacement reaction
The reaction will occur if one of the products is a solid. If one product is insoluble then the reaction will occur.
The reaction will occur if one of the products is a gas.
The reaction will occur if one of the products is water.