Chemical Equations. What are they? Equations showing chemical change(s). Example: CH 4 + 2O 2 CO 2 +...
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Chemical Equations. What are they? Equations showing chemical change(s). Example: CH 4 + 2O 2 CO 2 + 2H 2 O **Law of Conservation of Mass: reactantsproducts
What are they? Equations showing chemical change(s). Example:
CH 4 + 2O 2 CO 2 + 2H 2 O **Law of Conservation of Mass:
reactantsproducts In a chemical reaction, mass (atoms) cannot be
created or destroyed.
Slide 3
A closer look Mass is always conserved in a chemical reaction.
CH 4 + 2O 2 CO 2 + 2H 2 O ReactantsProducts Carbon11 Hydrogen44
Oxygen44
Slide 4
Most chemical equations give the physical states of the
reactants and products: (s) solid (l) liquid (g) gas (aq) aqueous
solution (dissolved in water) Ex: 2Mg(s) + O 2 (g) 2MgO(s) 2HgO(s)
2Hg(l) + O 2 (g)
Slide 5
Common diatomic gases H 2 N 2 O 2 Cl 2 Br 2 I 2 F 2
Slide 6
6 Common Types of Chemical Reactions: 1) Synthesis 2)
Decomposition 3) Combustion 4) Single Replacement 5) Double
Replacement 6) Oxidation Reduction (REDOX) Reaction
Slide 7
Synthesis Reaction in which 2 or more reactants combine to form
a new compound. A + B AB Examples: 2K(s) + Cl 2 (g) 2KCl(s) Fe(s) +
S(s) FeS(s)
Slide 8
Decomposition A single compound is broken down into 2 or more
products AB A + B Examples: CaCO 3 (s) CaO(s) + CO 2 (g) 2H 2 O(l)
2H 2 (g) + O 2 (g) PbO 2 (s) Pb(l) + O 2 (g) 2HgO(s) 2Hg(l) + O
2
Slide 9
Combustion (burning in presence of Oxygen) Oxygen gas (O 2 )
reacts with another substance to produce an oxide*, water, and
heat. *The oxide is typically carbon dioxide. Examples: CH 4 (g) +
2 O 2 (g) CO 2 (g) + 2 H 2 O(g) + heat C 3 H 8 (g) + 5 O 2 (g) 3 CO
2 (g) + 4 H 2 O(g) + heat
Slide 10
Single Replacement One element replaces another element in a
compound A + BC AC + B (a metal replaces a cation) AB + C AC + B (a
nonmetal replaces an anion) Examples: Mg(s) + Zn(NO 3 ) 2 (aq)
Mg(NO 3 ) 2 (aq) + Zn(s) 2K(s) + 2H 2 O(l) 2KOH(aq) + H 2 (g)
Slide 11
You must follow the rules! NO REACTION
Slide 12
The Activity Series of Metals A list of metals in order of
decreasing reactivity. A reactive metal will replace any of the
metals listed below it in the series. Examples: Mg(s) + AgNO 3 (aq)
Mg(NO 3 ) 2 (aq) + Ag(s) Mg(s) + LiNO 3 (aq) No Reaction
Slide 13
Halogen Activity Series
Slide 14
Double Replacement An exchange of cations between two reacting
compounds. This type of reaction typically occurs between 2 ionic
compounds in aqueous solution. AB + CD AD + CB Example: AgNO 3 (aq)
+ KCl(aq) AgCl(s) + KNO 3 (aq)
Slide 15
3 Types of Double Replacement Reactions 1. Precipitation
Reaction -produces a precipitate (formation of a solid) -Example: K
2 CO 3 (aq) + BaCl 2 (aq) 2KCl(aq) + BaCO 3 (s) 2. Gas Production
-produces a gas -Example: 2NaCN(aq) + H 2 SO 4 (aq) 2HCN(g) + Na 2
SO 4 (aq) 3. Neutralization Reaction -produces water and salt
-Example: Ca(OH) 2 (aq) + 2HCl(aq) CaCl 2 (aq) + 2H 2 O(l)
Slide 16
Oxidation Reduction (REDOX) Reactions Any reaction in which the
oxidation states of atoms change Oxidation=loss of electron(s)
Reduction=gain of electron(s) Example: 4Al (s) + 3O2 (g) 2Al2O3 (s)
4Al 4Al + 12 e - Al was oxidized 6O + 12 e - 6O O was reduced 00 +3
0 -2 0 +3 -2
Slide 17
Assigning Oxidation States (aka Oxidation Number) Hypothetical
charge use to indicate the degree of oxidation (loss of electrons)
Rules in assigning oxidation states: 1) The oxidation state of a
free element is zero (0). ex. O 2 (g), Ag (s) 2) The oxidation
state of a monatomic ion is equal to its ionic charge. (ex. Na +,
Cl -3 ) 3) H has an oxidation of +1 and oxygen has an oxidation
state of -2 when they are present in most compounds. (Exceptions:
Hydrogen is -1 in hydrides of metals such as LiH, and Oxygen is -1
in peroxides such as H 2 O 2.) 4) The sum of all oxidation states
in a neutral molecule must be zero, and the sum of the oxidation
states in an ion must be equal to the ions overall charge.
Slide 18
Writing and Balancing Chemical Equations: Step 1) Determine the
reactants, products and the physical states. Step 2) Write the
unbalanced equation to summarize the word equation. Step 3) Use
coefficients to balance each and every element EXCEPT H and O. Step
4) Use coefficients to balance the Hydrogens. Step 5) Use
coefficients to balance the Oxygens.
Slide 19
Example #1 Chlorine gas and potassium iodide solution react
together to form potassium chloride and iodine gas. Cl 2 (g) +
KI(aq) KCl(s) +I 2 (g) Cl 2 (g) + 2KI(aq) 2KCl(s) +I 2 (g) Write
the equation: Balance the equation:
Slide 20
Example #2 Aluminum hydroxide solution is heated and decomposed
into a solid aluminum oxide and liquid water. Write the
equation:Al(OH) 3 (aq) Al 2 O 3 (s) + H 2 O(l) Balance the
equation2Al(OH) 3 (aq) Al 2 O 3 (s) + 3H 2 O(l)
Slide 21
Example #3 Zinc and Lead (II) nitrate react to form zinc
nitrate and Lead. Zn(s) + Pb(NO 3 ) 2 (s) Zn(NO 3 ) 2 (s) +
Pb(s)
Slide 22
Predicting Products while considering the Solubility Rules:
Example: K 2 CrO 4 (s) + Ba(NO 3 ) 2 (s) KNO 3 (s) + BaCrO 4
(s)