Chemical Equations and Reactions Mrs. Partridge Loveland HS Chemistry Chemical Equations and Reactions Mrs. Partridge Loveland HS

Representing Chemical Changes A chemical reaction is the changing of substances to other substances by the breaking of bonds in reactants and the formation of bonds in products. Chemical reactions are represented by chemical equations which show what changes take place and the relative amounts of the various elements and compounds that take part in these changes.

Evidence for Chemical Reactions A gas is produced A precipitate (solid) is formed when 2 solutions are combined A permanent color change occurs Heat energy change is noted

The Logistics of a Chemical Equation Every thing to the left of the arrow is considered a reactant. These are the starting substances. Everything to the right of the arrow is considered a product. These are the substances that are formed as a result of the chemical reaction. Coefficients go before the chemical formula and tell how many molecules, atoms or formula units of that species react or are formed.

The Logistics of a Chemical Equation A double arrow indicates that the reaction is reversible. A small triangle Δ over the reaction arrow means that heat was added to make the reaction occur.

The Logistics of a Chemical Equation The small letters in parenthesis indicate the state of matter of the reactants and products H2O(l )  2H2(g) + O2(g) (s) = solid (l) = liquid (g) = gas (aq) = aqueous  dissolved in water

The Logistics of a Chemical Equation Arrows can also be used to indicate state of matter on the product side. Up arrow = gas Down arrow = precipitate

Reading a Chemical Equation Read the equation by naming the state and the name of the reactants first. The arrow is read as “yields” or “produces”. The products are then read with the state and name.

Reading a Chemical Equation EXAMPLE: 2Al (s) + Fe2O3 (s)  2Fe (s) + Al2O3 (s) Solid aluminum reacts with solid iron (III) oxide to yield solid iron and solid aluminum oxide

WHEN GOING FROM THE WORD EQUATION TO THE CHEMICAL EQUATION, MAKE SURE TO PAY ATTENTION TO DIATOMIC MOLECULES!

Balancing Equations The purpose of balancing equations is to ensure that the law of conservation of mass is maintained. You must show an equal number of atoms for each element on both sides of the equation. Remember, atoms can not be created or destroyed in a chemical reaction, they are ONLY rearranged!

Balancing Equations To balance an equation, NEVER touch the subscripts. Add coefficients in front of chemical formulas until there are equal number of each type of atom on the reactant and the product side.

Balancing by Inspection Step One: Write down what type of elements are present in the chemical equation Step Two: Count the number of elements on the reactant side and on the product side Step Three: Add coefficients to get the same number of elements on both sides of the equation

Balancing by Inspection HINTS: Balance compounds first Treat polyatomics as one unit if they travel (stay) together Leave any element that stands alone for LAST Carbon and hydrogen are usually best left for last

Balancing by Algebra Step One: Put a lower case letter in front of every chemical formula Step Two: Write an equality that relates the reactant side to the product side for each type of element Step Three: Assign the lower case letter with the highest coefficient the value of one Step Four: Solve algebraically to determine the remaining letters Step Five: The final answers must be WHOLE NUMBERS

Types of Chemical Reactions Synthesis or Composition Reaction Decomposition Reaction Single Replacement Reaction Double Replacement Reaction Combustion Reaction

Synthesis (Composition) Reactions Two or more substances combine to form a new compound. A + X  AX Reaction of elements with oxygen and sulfur Reactions of metals with Halogens Synthesis Reactions with Oxides There are others not covered here!

Decomposition Reactions A single compound undergoes a reaction that produces two or more simpler substances AX  A + X Decomposition of: Binary compounds H2O(l )  2H2(g) + O2(g) Metal carbonates CaCO3(s)  CaO(s) + CO2(g) Metal hydroxides Ca(OH)2(s)  CaO(s) + H2O(g) Metal chlorates 2KClO3(s)  2KCl(s) + 3O2(g) Oxyacids H2CO3(aq)  CO2(g) + H2O(l )

Single Replacement Reactions A + BX  AX + B BX + Y  BY + X Replacement of: Metals by another metal Hydrogen in water by a metal Hydrogen in an acid by a metal Halogens by more active halogens

The Activity Series of the Metals Lithium Potassium Calcium Sodium Magnesium Aluminum Zinc Chromium Iron Nickel Lead Hydrogen Bismuth Copper Mercury Silver Platinum Gold Metals can replace other metals provided that they are above the metal that they are trying to replace. Metals above hydrogen can replace hydrogen in acids. Metals from sodium upward can replace hydrogen in water

The Activity Series of the Halogens Fluorine Chlorine Bromine Iodine Halogens can replace other halogens in compounds, provided that they are above the halogen that they are trying to replace. 2NaCl(s) + F2(g)  2NaF(s) + Cl2(g) ??? MgCl2(s) + Br2(g)  No Reaction ???

Double Replacement Reactions The ions of two compounds exchange places in an aqueous solution to form two new compounds. AX + BY  AY + BX One of the compounds formed is usually a precipitate, an insoluble gas that bubbles out of solution, or a molecular compound, usually water.

Combustion Reactions A substance combines with oxygen, releasing a large amount of energy in the form of light and heat. Reactive elements combine with oxygen P4(s) + 5O2(g)  P4O10(s) (This is also a synthesis reaction) The burning of natural gas, wood, gasoline C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)

Net Ionic Equations Double replacement reactions occur when 2 ionic compounds react by exchanging cations to form 2 different compounds.

Net Ionic Equations When 2 solutions of ionic compounds are mixed the formation of a precipitate, gas or water may be one of the products. One of the products is only slightly soluble and precipitates form solution Na2S(aq) + Cd(NO3)2  CdS(s) + 2NaNO3 (aq)

Net Ionic Equations One of the products is a gas that escapes out of the mixture. 2NaCN(aq)+H2SO4(aq)2HCN(g)+Na2SO4(aq) One of the products is a molecular compound such as H2O. Ca(OH)2(aq)+ 2HCl  CaCl2(aq) +2H2O(L)

To predict the formation of a precipitate we use SOLUBILITY RULES! Net Ionic Equations To predict the formation of a precipitate we use SOLUBILITY RULES! **See appendix in book.

Net Ionic Equations The world is water based. More than 70% of the earth’s surface is covered with water. Consequently, many important chemical reactions take place in water  aqueous solutions.

Net Ionic Equations AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq) Most ionic compounds dissociate – separate into anions and cations when they dissolve in water. A complete IONIC EQUATION shows dissolved ionic compounds as their free ions. Ag+(aq)+ NO3-(aq) + Na+(aq) + Cl-(aq)  AgCl(s) + Na+(aq) + NO3-(aq)

Net Ionic Equations This equation can be simplified by eliminating the IONS that don’t participate in the reaction, spectator ions, from both sides of the equation. Ag+(aq)+NO3-(aq)+Na+(aq)+Cl-(aq)AgCl(s) + Na+(aq)+NO3-(aq) Rewritten, this equation is called the NET IONIC EQUATION, indicating only those particles that actually take part in the reaction. Ag+(aq) + Cl-(aq)  AgCl(s) Net ionic equations must have balanced ionic charges.