Chemical Formulas and equations

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

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

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

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

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.

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

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.

Writing formulas and naming compounds

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

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.

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

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

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

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?

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

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)

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

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

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.

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.

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)

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

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

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

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

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

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.