1. A chemical reaction involves the breaking and forming of chemical bonds between atoms or ions, resulting in new substances. 2. Evidence of a chemical reaction includes color changes, formation of a precipitate or gas, and temperature changes. 3. A chemical equation represents reactants yielding products, with coefficients indicating relative amounts and state symbols denoting physical states. Balancing equations ensures the same number and type of atoms are on both sides.

1. EQUATIONS &
REACTIONS

2. A. Chemical Changes and Reactions
1. New substances are producedproduced.
2. Chemical reaction – chemical bonds between atoms or ions breakbreak, and new bondsnew bonds form
between atoms or ions.
B. Evidence of a Chemical Reaction
1. color changecolor change
2. formation of a precipitateformation of a precipitate
3. temperature changetemperature change
4. formation of a gasformation of a gas
1. Describing Reactions

3. Describing Chemical ReactionsDescribing Chemical Reactions
• Cellular phone messages make use of
symbols and abbreviations to express
ideas in shorter form. Similarly, chemists
often use chemical equations in place of
words.

4. C. Mechanics of a Chemical Reaction
1. Starting Materials – reactantsreactants
2. Ending Materials - productsproducts
3. reactantsreactants → productsproducts
Arrow = yieldsyields or producesproduces
4. If there are multiple products or reactants, they are connected with a +
symbol.
5. Symbols above the yield sign represent conditionsconditions necessary for a
reaction to proceed.
Ex)
= delta = heatheat
= electrolysiselectrolysis
Δ→
elec→

5. 5. Some reactions occur spontaneouslyspontaneously.
6. Symbols represent the statestate of the reactants and products.
Liquid = ll Gas = gg Solid = ss Crystal = crcr
Aqueous = aqaq (solids in water solutionwater solution)
DEMO or video https://www.youtube.com/watch?v=Hv15HxOVpzE
Ex) 2Al(s) + 3CuCl2(aq) → 2AlCl3(aq) + 3Cu(s)
silver blue gray red

6. 7. Complete chemical equations include the subscript to
indicate the physical statestate of each substance.
8. Diatomic molecules – certain elements exist in
nature as diatomic molecules (X2)
List them: N2 O2 F2 Cl2 Br2 I2 H2

7. Natural States of the Elements
■ Diatomic Molecules
Nitrogen gas contains
N2 molecules.
Oxygen gas contains
O2 molecules.

8. Write a Word Equation of
Chemical Reacting
• Put the name of the reactants on the left,
separating with a +.
• Use an arrow to replace the verb for
reacting.
• Put the names of the products on the
right, separated by + signs

9. Solid calcium and water react to
yield calcium hydroxide and
hydrogen gas.

10. Solid calcium and water react to
yield calcium hydroxide and
hydrogen gas.
Ca(s) + H2O -> Ca(OH)2 + H2

11. Write a Skeleton Equation of
Chemical Reactions
• The symbols for reactants on the left,
separating with a +.
• An arrow to replace the verb for reacting.
• Products formulas are on the right,
seperated by + signs.
• If You are starting with a word equations,
you are just replacing the names with the
formulas.

12. Sodium hydroxide and
hydrochloric acid produce
sodium chloride and water.

13. Sodium hydroxide and
hydrochloric acid produce
sodium chloride and water.

14. Is it Balanced ?
■ A balanced equation has the same number of each type of
molecules in the reactants and products.

15. Balancing Chemical Equations
1. Conservation of massmass leads to balancing equations – the
number of atoms of each element must be the same
before & after the reactionbefore & after the reaction.
2. The Law of Conservation of Mass also states that the total
massmass before and after the reaction must be the samesame.
You cannot lose or gain mass.
3. Therefore the MASS OF THE PRODUCTS =
MASS OF REACTANTSMASS OF REACTANTS
4. Subscript – indicates number of atomsatoms of an element
present in a compound.
5. Coefficient – indicates the number of atomsatoms or
moleculesmolecules involved in the reaction.

16. 6. Steps to Balance Equations:
A. Write equation with symbols.
B. Count # of atoms on each side of the
reaction.
C. Balance atoms using coefficients.
D. General Rule: Balance all elements first.
Then, balance C, H, and O.
E. NEVER CHANGE SUBSCRIPTS!!!!

17. _____ Na + ____ H2O  ____ NaOH + _____ H2

18. 2 Na + 2 H2O  2 NaOH + H2

19. _____ NH3 + ____ NO  ___ N2 + ___ H2O

20. 4 NH3 + 6 NO  5 N2 + 6 H2O

21. 3 H2 + N2 → 2 NH3
3 Na2SO4 + Ca3(PO4)2 → 2 CaSO4 + 2 Na3PO4
2 NaNO3 → 2 NaNO2 + O2
2 C8H18 + 25 O2 → 16 CO2 + 18 H2O

22. DO NOW: Demonstration
When electricity is run through a container of
distilled water, bubbles form at both
electrodes. The negative end makes lots of
tiny bubbles really quickly. If the gases are
captured in vials, there is twice the amount
of gas at the negative end, than at the
positive end

23. A. Synthesis Reactions (direct combinationdirect combination)
1. Two or more elements or compoundselements or compounds combine
to form a more complexcomplex product.
A + B → AB
2. Ex. Fe + S → FeS
CaO + H2O → Ca(OH)2
2. Classifying Chemical Reactions

24. B. Decomposition Reactions (analysisanalysis)
1. A single reactantsingle reactant breaks down into simplersimpler
compounds or elements.
AB → A + B
2. The oppositeopposite of a synthesis reaction.
3. Ex. 2 HgO → 2 Hg + O2
CaCO3 → CaO + CO2

25. C. Single Replacement Reactions
1. Atoms of an uncombined element replacereplace atoms of
another element in a compound.
A + BX → AX + B
2. A moremore active element will replace a less active
element.
3. Ex: CuCl2 + Zn → ZnCl2 + Cu
Br2 +2 KI  2KBr + I2

26. D. Double-Replacement Reactions
1. Atoms or ions from two differentdifferent compounds
replace each other.
AX + BY → AY + BX
2. These types of reactions will
(A) form precipitatesprecipitates (↓)
(B) form gasesgases (↑)
(C) are acid-baseacid-base neutralizations
(produce water and salt)

27. Double Replacement Examples:
A. Pb(NO3)2 + 2 KI → PbI2 ↓ + 2 KNO3
B. CaCO3 + 2 HCl → CaCl2 + H2CO3
C. NaOH + HCl → NaCl + HOH
3. In letter B above, carbonic acid, H2CO3, is unstable
and will immediately decompose into carboncarbon
dioxidedioxide and waterwater.
CaCO3 + 2 HCl → CaCl2 + CO2↑ + H2O

28. E. Combustion Reactions
1. One substance reacts with oxygen,oxygen, OO22 to
produce oxide compounds.
2. Occurs during burningburning or oxidationoxidation
(rusting.)
3. The reactions that only add oxygen are
classified as synthesissynthesis reactions.
Ex) S + O2 → SO2

29. 4. Combustion reactions are exothermicexothermic, releasing a large
amount of energy as light, heat, or sound.
5. A true Combustion reaction occurs when a hydrocarbonhydrocarbon
(compound containing HH & CC ) reacts to form
carbon dioxide and waterwater that are always the
products.
CxHx + O2 → CO2 + H2O
6. Ex. __CH4 + __O2 → __CO2 + __H2O + 803 kJ
C6H12 O6 + O2 → CO2 + H2O + heat
22
666

33. Men with Hats

34. What kind of Reaction?

35. What kind of reaction?

38. CuCO3(s) → CO2(g) + CuO(s)
copper (II) carbonate
metal oxide
flame goes out

39. What is one of the products?
mercury (II) oxide
mercury (II) oxide
2HgO → 2Hg + O2

40. 3CuCl2 + 2Al → 2AlCl3 + 3Cu

41. Sodium Metal plus Chlori
2 Na + Cl2 → 2 NaCl
https://youtu.be/VBReOjo3ri8

42. 3. The Mole

43. Percent Composition is the percent of any
element contributes to the mass of a
compound.
4. Chemical Composition

44. Steps to Calculate % Comp
1. Write formula
2. Find mass of individual elements
3. Total for molecular (formula) mass.
4. Part/whole x 100 = %
• Can be used to find the mass of an
element in compound or solution.

45. Potassium Permanganate

46. Empirical Formula
• Empirical- Relying on or derived from
observation or experiment. Verifiable or
provable by means of observation or
experiment.

47. Empirical Formula
• The empirical formula is the simplest
formula for a compound. A molecular
formula is the same as or a multiple of the
empirical formula, and is based on the
actual number of atoms of each type in
the compound. For example, if the
empirical formula of a compound is C3H8 ,
its molecular formula may be C3H8 , C6H16 ,
etc.

48. Empirical?
• H2O
• H2O2
• Mg(OH)2
• C10H14O
• C6H12O6
• Hg2 (CN)2

49. Steps to find Empirical Formula
1. Find grams of each element (assume
100g)
2. Change grams to moles
3. Look for smallest whole number ratio
(divide by smallest)
• Compare to mass molecular mass to find
molecular formula