1. Heat Energy and Chemical Equations Part 1: Changes in Matter & Energy Balancing Equations Types of Reactions

2. Physical vs Chemical Changes in Matter & Energy
Matter is constantly undergoing changes. These changes can be identified as physical or chemical
Physical Changes is any change that does not result in new substances formed. The original substance is just in a new form.
For Example
Phase Changes and Dissolving
H2O(s) H2O(l)
NaCl(s) Na+1(aq) + Cl-1(aq)

3. Changes in Matter (continued)
Chemical Changes occur when substances collide and change into new substances.
This occurs because the bonds of the substances you start with break and new bonds are formed resulting in new substances.
When a chemical change occurs it is called a chemical reaction.
For Example
Burning of paper & Rusting of a metals

4. Chemical Equations
A chemical equation can be used to show the changes that take place during a chemical or physical change.
The substances that enter the reaction are written on the left side of the equation and called reactants.
If the reaction involves more than one substance they will be separated by a “+” sign.
The new substances that are formed from the reaction are known as products and they are written on the right side of the equation.
Separating the reactants and the products is an arrow and is read as yields or produces.
For Example
H2O(s) + energy H2O(l)
CH4 + O CO2 + 2H2O + energy

5. Energy and Changes in Matter
Energy is defined as the ability to do work and it is NOT matter
Energy is measured in the unit Joules (J) See Table D
Can be POTENTIAL or KINETIC

6. Potential Energy Energy that is stored (i.e. in a chemical bond).
Something has the “potential” to do some kind of work
Example: the child at the top of the slide has potential energy

7. Kinetic Energy Energy of motion
Example: the child going down the slide now has kinetic energy

8. Law of Conservation of Energy
Energy, like matter, is neither created nor destroyed, rather it is converted.

9. Endothermic and Exothermic
Chemical and Physical changes always involve the loss and gain of energy.
This energy is most often expressed or described as heat.
Based on whether energy is absorbed or released you can classify energy changes as either endothermic or exothermic.

10. Understanding Heat Flow
Heat is defined as the energy that transfers from one object to another.
Heat will always flow from warm  cool.
What will happen if the two objects are touching?

11. Heat Energy vs. Temperature
We can measure heat flow by identifying changes in temperature.
When an objects temperature goes down it is losing heat when it goes up it is absorbing heat.
Temperature is a measure of the average kinetic energy or speed of the particles in matter.

12. Heat Energy and Changes in Matter
In virtually all changes in matter, energy is released or absorbed.
System vs. Surroundings
The system is what ever is being observed the surroundings are everything around the system.
In a chemical change we can measure changes of the surroundings to identify whether or not the reaction (the system) gained or lost energy.

13. Examples

14. Exothermic Processes (Changes)
Exothermic processes RELEASE ENERGY (i.e. explosions).
A good way to remember this is to associate “EXO” with “OUT”.
They have a –ΔH value because heat is leaving the system.
The surrounding temperature increases because the system or reaction in this loses energy
Heat is a product and will be written on the right side of the equation.

15. Endothermic Processes (Changes)
Endothermic processes ABSORB ENERGY
A good way to remember this is to associate “ENDO” with “INSIDE”.
Has a +ΔH value because heating is entering the system.
The surrounding temperature decreases because the system or reaction absorbs heat from the surroundings make it feel cold.
Heat is a reactant written on the left side of the equation.

16. Examples of Endothermic and Exothermic Processes in the form of equations

17. ** mass must be conserved
Balancing Equations
Based on the Law of Conservation of Mass
In a closed system, the mass of the reactants must equal the mass of the products.
Example: formation of water
H O2  H2O
** mass must be conserved

18. H2 + O2  H2O 2g + 32g = 18g 2H2 + O2  2H2O 4g + 32g = 36g
How do we get these masses to balance?
2H O2  2H2O
4g g = g

19. 2H2O

20. Counting atoms when balancing:
The number of atoms of each element must be equal on each side of the reaction. Therefore the masses will be equivalent.
2H O2  2H2O
4H 2O = 4H 2O

21. Rules to Balance Equations:
1. Start with the element that is only found once on both sides.
2. Keep polyatomic ions together. Count as a unit if not broken up.
3. Coefficients must be smallest possible whole number.

22. Al2 (SO4) 3 + CaCl2  AlCl3 + CaSO4
Examples:
K Cl2  KCl
Al2 (SO4) CaCl2  AlCl3 + CaSO4
H2O + CO2  H2CO3 2 2 3 2 3

23. Try These:
H I2  HI
Na + Cl2  NaCl
Al + O2  Al2O3 2 2 2 4 2 3 2

24. Phases in equations
You must write in the phase of each reactant and product.
(s) Solid Mg (s) magnesium metal
(l) Liquid Br2 (l) bromine liquid
(g) Gas CH4 (g) methane gas
(aq) aqueous NaBr (aq) sodium bromide dissolved in water

25. Summary
Chemical equations are written to show the chemical change from reactants to products.
Endothermic Reactions absorb heat. Heat is a reactant.
Exothermic reactions release heat. Heat is a product.
Mass of reactants = Mass of products
Equations must be balanced so that the number of atoms of each element on each side of the equation are equal.

26. Types of Chemical Reactions
There are 4 main kinds of reactions that occur between elements and compounds.
Each of these chemical reactions follow a specific format from reactants to products. If you recognize the format you will be able identify the type of reaction.

27. SYNTHESIS REACTIONS Also called combination reactions.
General Equation Format
A + B  AB
***These reactions will only have ONE product***
Example: 2H2 + O2  2H2O
This reaction represents the “Synthesis of Water”

28. DECOMPOSITION REACTIONS
Also called analysis reactions.
General Equation Format
AB  A + B
***You can recognize this because it only has ONE reactant***
Example: CH4  C + 2H2
This reaction is known as the “Decomposition of Methane”

29. SINGLE REPLACEMENT
Often yield such gases as H2 when it involves an acid.
AB + C  AC + B
*** Notice these reactions have a compound and an element as a reactant and product***
Example: Mg + 2HCl  MgCl2 + H2

30. DOUBLE REPLACEMENT AB + CD  AD + CB
Two elements in two different compounds switch or replace one another.
Often occurs in aqueous solutions of ionic compounds.
AB + CD  AD + CB
*** These reactions will have two compounds on the reactant and product sides.***
Example: 3NaCl + AlPO4  Na3PO4 + AlCl3

31. Summary- There are 4 Basic Types of Reactions
Synthesis
Decomposition
Single replacement
Double replacement

32. Chemical Equations Part 2
Predicting Products of Reactions
Writing Equations

33. Predicting Products of Reactions
How do we predict the products of a reaction?
First you must know the types of reactions.
Fill in the types on your notes.

34. Types of Reactions
Synthesis – 2 or more reactants 1 product Decomposition – 1 reactant 2 or more products SRR – element + compound compound + element DRR – switch positive ions of reactants to form products

35. Single Replacement Reactions
In a single replacement reaction the lone element will replace the element in compound that has the same ionic charge.
Zn + HCl >>>> “Zn” forms a positive ion so it will replace“H” the positive ion in HCl.
Br2 + HCl >>>>> “Br” forms a negative ion so it will replace “Cl” the negative ion in HCl.
Do they always occur?
No!!!!!!!!!!!
You can predict if a SRR occurs by studying the elements that are switching.

36. SRR Example: Zn + 2HCl  ZnCl2 + H2
If the element lone element is more reactive (active) than the element it would replace then the reaction will occur (spontaneous = reaction occurs) if the element is less reactive then the reaction will not occur (nonspontaneous = does not occur)
In the reaction above Zinc is more reactive so the Zinc will replace the Hydrogen in the compound to form Zinc Chloride.
The Hydrogen would leave the reaction as a gas.

37. How do we know if one element is more reactive than another?
Look at Table J on the reference tables.
Table J is an Activity Series.
This lists the reactivity of metals and non-metals.
NOTE: Ag and Au are the least reactive metals. F2 is the most reactive non-metal.

38. How do we use the Activity Series?
In a SRR, if the element by itself is more reactive than the element it would replace than the reaction will occur or is spontaneous.
Ex: Zn + HCl  will occur

39. If a reaction occurs… Switch and write the products. Balance.
Zn + HCl  ZnCl2 + H2 2

40. What if the reaction does not occur?
Write “ no reaction”.
No need to balance
Example:
Mg + Li NO3  no reaction

41. Try these… Mg + Zn(NO3)2  Mg + AgNO3  Cl2 + NaBr  Sn + NaNO3 
Br2 + KF 
Cu + AgNO3 
MgCl2 + Na 
Mg(NO3)2 + Zn 2
Mg(NO3)2 + Ag 2 2
NaCl + Br2 2
No reaction
No reaction
CuNO3 + Ag 2
NaCl + Mg 2

42. Double Replacement Reactions
All compounds are ionic.
Products are formed by switching the positive ions of the reactants.
To write products write the ions and use the criss-cross rule.

43. Ionic Compounds in Water
Break up into ions
Example 1: NaCl in water becomes Na+ + Cl-
Example 2: K3PO4 in water becomes K+ and PO4-3

44. How do we know if a DDR occurs?
DDR “go to completion” or occur if a solid, liquid, or gas forms:
(s) = useTable F
(l) = H2O
(g) = H2, Cl2, etc.
**If all products are (aq) then no reaction happens.

45. Using Table F to Predict a Solid
DDR occurs if one product formed is insoluble (solid that precipitates out of solution).
Check Table F for Solubility Guidelines.
Soluble = no precipitate, no rxn (aq)
Insoluble = precipitate, rxn occurs (s)
If both products are soluble (aq) then reaction does not occur.

46. DDR Example: Na2CO3 (aq) + Ca(OH)2 (aq) 
Write the ions above each reactant
Switch the positive ions and write new ion pairs on product side (don’t forget “+ ion” first
Write the formulas for the products using the criss-cross rule.
Check products with Table F guidelines.
NaOH is soluble so it is labeled (aq) it dissolves (see Group 1 ions)
CaCO3 is insoluble (s) it will not dissolve forms a percipitate. (see carbonate)
5) Balance the equation
Na+1 CO Ca+2 OH-1
Ca+2 CO Na+1 OH-1 2
Na2CO3 (aq) + Ca(OH)2 (aq) 
CaCO3 (s) + NaOH (aq)

47. Try these… Write products and predict if it “goes to completion”
NH CO Ca+2 Cl-1
(NH4)2CO3 (aq) + CaCl2(aq) 
K+1 NO Ca+2 I-1
KNO3 (aq) + CaI2 (aq) 
Na+1 OH H+1 SO4-2
NaOH(aq) + H2SO4 (aq) 
Cu+2 SO Na+1 OH-1
CuSO4 (aq) + NaOH (aq) 

48. Writing the Products (NH4)2CO3 (aq) + CaCl2(aq) 
NH CO Ca+2 Cl-1
(NH4)2CO3 (aq) + CaCl2(aq) 
K+1 NO Ca+2 I-1
KNO3 (aq) CaI2 (aq) 
Na+1 OH H+1 SO4-2
NaOH(aq) H2SO4 (aq) 
Cu+2 SO Na+1 OH-1
CuSO4 (aq) + NaOH (aq) 
NH4+1 Cl Ca+2 CO3-2
NH4Cl (aq) CaCO3(s)
K+1 I Ca+2 NO3-1
KI (aq) Ca(NO3)2 (aq)
Na+1 SO H+1 OH-1
Na2SO4 (aq) H2O (l)
Cu+2 OH Na+1 SO4-2
Cu(OH)2 (s) Na2(SO4) (aq)

49. Summary
Use Activity Series (Table J) to predict the products of single replacement reactions.
Use Solubility Guidelines (Table F) to predict the products of double replacement reactions.

50. Predicting Missing Reactant or Product
Based on law of conservation of mass matter can neither be created or destroyed
When given a balanced equation you should be able to determine the formula of missing substance
Count the atoms on both sides, subtract the atoms on the missing formula side from the side with the known formulas
Any missing element must be present in the unknown formula

51. 2Na + 2H2O  x + 2NaOH Na = 2 Na = 2 H = 4 H = 2 O = 2 O = 2
Whats missing?
2 Hydrogens
X must be H2

52. Try These 1) NaO 2) NaOH 3) Na2O 4) Na2OH 4Fe + 3O2  2X 1) FeO
Multiple Choice Questions
2Na H2O  2x + H2
1) NaO
2) NaOH
3) Na2O
4) Na2OH
4Fe + 3O2  2X
1) FeO
2) Fe2O3
3) Fe3O2
4) Fe3O4