This document provides an overview of chemical reactions and special topics covered in Unit 4. It begins by describing how to write chemical equations, including identifying reactants and products, and balancing equations. It then discusses the four main types of chemical reactions: synthesis, decomposition, single replacement, and double replacement. The document also covers topics like energy changes in reactions, explaining that energy is either absorbed or released. It provides examples of exothermic and endothermic reactions. Finally, it discusses solutions and solubility, radioactive decay processes, and half-life.

1. UNIT 4:
CHEMICAL REACTIONS & SPECIAL
TOPICS

2. BILL NYE THE SCIENCE GUY
 http://www.youtube.com/watch?v=PlwuxpMh8nk

3. 7.1: DESCRIBING REACTIONS
The most useful way of describing reactions (or a
chemical change) is by stating what was present
before and after the chemical change.

4. CHEMICAL CHANGE
 chemical reaction
 process in which one or more substances are changed into new
substances
 reactants
 substance that reacts
 product
the new substance that is formed

5. CONSERVATION OF MASS
The French chemist Antoine Lavoisier
established that the total mass of the products
always equals the total mass of the reactants.
 This is called the Law of Conservation of Mass (mass is
not created nor destroyed in a chemical reaction).

6. WRITING EQUATIONS
chemical equation
shorthand method to describe a chemical
reaction using chemical formulas and
other symbols
Reactants
HgO(s) 
Products
Hg(l) + O2(g)

7. WRITING EQUATIONS
(aq) – aqueous (substance dissolved in water)
(s)– solid
(l) – liquid
(g) – gas
coefficients – the numbers to the left of the formulas used to help
balance the equation

8. BALANCING EQUATIONS
 If you notice that the number of atoms on the left side does
not equal the number of atoms on the right side then we
must balance the equation.
 Since mass is conserved before and after a chemical reaction, an equation MUST BE
balanced in order for it to be true.

9. BALANCING EQUATION RULES
1. You should NEVER change the subscripts in a
formula.
2. Start by, counting the number of atoms of each
element on each side of the equation.
3. Change one or more coefficients until the
equation is balanced

10. EXAMPLE
 Balance the following equation
NiCl2(aq) + NaOH(aq)  Ni(OH)2(s) + NaCl(aq)

11. EXAMPLE
HgO(s)  Hg(l) + O2(g)

12. 7.2: TYPES OF REACTIONS
 Reactions
are classified by the type of reactant or
the number of reactants and products.
There are 4 different types of reactions that we
will discuss
 Synthesis
 Decomposition
 Single-Replacement
 Double-Replacement

13. SYNTHESIS
A synthesis reaction is a reaction in which two or
more substances react to form a single substance.
 The product synthesized is always a compound
 Examples
A + B  AB
2Na + Cl2 2NaCl

14. DECOMPOSITION
 The opposite of synthesis
A decomposition reaction is a reaction in which a
compound breaks down into two or more simpler
substances.
 The reactant MUST BE a compound.
 Examples
AB A + B
2H2O 2H2 + O2

15. SINGLE REPLACEMENT
A single replacement reaction is a reaction in
which one element takes the place of another
element in a compound.
Example Form:
A + BC B + AC
Cu + 2AgNO3  2Ag + Cu(NO3)2

16. DOUBLE REPLACEMENT
A double replacement reaction is one in which
two different compounds exchange positive ions
and form two new compounds.
Example Forms:
AB + CD AD + CB
Pb(NO3)2 + 2KI PbI2 + 2KNO3

17. 7.3: ENERGY CHANGES IN REACTIONS
Chemical Energy – the energy stored in the
chemical bonds of a substance.
Chemical Reactions involve the breaking of
chemical bonds in the reactants and the formation
of chemical bonds in the products.

18. BREAKING BONDS
Breaking Bonds REQUIRES energy.
 This means we need to ADD energy in order to break
the bonds of reacting molecules in order to get the
reaction started.

19. FORMING BONDS
The formation of chemical bonds RELEASES
energy.
 When new chemical bonds are formed, a bit of energy is
released usually in the form of heat or light.

20. ENERGY IN REACTIONS
During a chemical reaction, energy is either
ABSORBED or RELEASED.
 We describe these reactions in two different ways either
Exothermic or Endothermic.

21. EXOTHERMIC REACTIONS
A chemical reaction that RELEASES energy is called
an exothermic reaction.
 The energy released as the products form is greater
than the energy required to break the bonds in the
reactants.
 Think of it as energy is EXITING the reaction
 EXiting _ Exothermic

22. ENDOTHERMIC REACTIONS
A chemical reaction that absorbs energy from its
surroundings is called an endothermic reaction.
 This means that there is more energy require to break
the bonds of the reactants than is released by the
formation of the products.

23. CONSERVATION OF ENERGY
The total amount of energy BEFORE a reaction is
EQUAL to the total amount of energy AFTER a
reaction.
 This is called the Conservation of Energy.

24. 8.1: FORMATION OF SOLUTIONS
A solution is a mixture that forms when
substances dissolve and form a homogeneous
mixture
 In order for a solution to form, one substance must
dissolve in another.

25. DISSOLVING
Every solution has two components
 A solute is a substance who particles are dissolved in a
solution
 A solvent is the substance in which the solute dissolves
in.
There are three ways that substances can dissolve
into water: dissociation, dispersion, ionization

26. DISSOCIATION
 In order for a solution to form, the attractions that hold the solute
together and the solvent together must be overcome.
The process in which an ionic compound separates
into ions as it dissolves is called dissociation.
 Example: Sodium Chloride & Water

27. DISPERSION
Sugar dissolves into water by dispersion, or
breaking into small pieces that spread throughout
the water.
 Example: Sugar & Water

28. IONIZATION
The process in which molecules gain or lose
electrons is known as ionization.
 Example: Ions are formed by the reaction of the solute
and solvent particles.

29. PROPERTIES OF LIQUID SOLUTIONS
Conductivity: ability to conduct electric current
Boiling Point: temperature needed for solution to
change from liquid phase to gas phase
Freezing Point: temperature needed for solution to
turn from liquid phase to solid phase.
Solutions can also be described as endothermic or
exothermic depending upon whether energy is
released or absorbed.

30. 8.2: SOLUBILITY & CONCENTRATION
The maximum amount of solute that dissolves in a
given amount of solvent at a constant temperature
is. called solubility
 Depending upon the amount of solute in a solution,
solutions can be described as either saturated,
unsaturated or supersaturated.

31. SATURATED
A saturated solution is one that contains as much
solute as the solvent can hold at a given
temperature.

32. UNSATURATED
A solution that has less than the maximum amount
of solute that can be dissolved is called an
unsaturated solution.

33. SUPERSATURATED
A supersaturated solution is one that contains
more solute than it can normally hold at a given
temperature.
 Usually very unstable.

34. FACTORS AFFECTING SOLUBILITY
Polarity of the solvent
 “like dissolves like”
 Solution formation is more likely to happen when the solute
and solvent are either both polar or both nonpolar.
Temperature
 The solubility of a solids increases as the solvent
temperature increases
Pressure
 Increasing pressure on a gas increases solubility in a liquid.

35. CONCENTRATION
The concentration of a solution is the amount
solute dissolved in a specified amount of solution.
 Can be expressed as percent by volume, percent by
mass and molarity.

36. SOLUBILITY CURVE
 Each line on the graph is
called a solubility curve for a
particular substance.
 You can use a solubility
curve to figure out how
much solute will dissolve at
any temperature given on
the graph.

37. 10.1: RADIOACTIVITY
Henri Becquerel
 1896
 left uranium salt in a drawer with a photographic
plate
 when he developed the plate, he found an outline
of the clumps of the uranium salt
 he hypothesized that the uranium salt emitted some
sort of energy
Marie and Pierre Curie
 students of Becquerel
 2 years later, they discovered Po and Ra while
studying uranium ore “pitch blende”

38. THEY DISCOVERED …
 Radiation
 release of matter and energy from nucleus
 Light energy (electromagnetic spectrum)
 all forms of radiation

39. STRONG FORCES
Protons are held together by strong
forces
short range force
as the distance increases, the force weakens
causes protons and neutrons to be attracted to
each other

40. STRONG FORCES
 to hold a nucleus together tightly, the
nucleus can decay and give off matter and
energy
 stable nucleus
 stays together permanently
 unstable nuclei
 radioactive!!!!
 nucleus does not stay together; emits matter and energy

41. RADIOACTIVE ELEMENTS
 radioactivity
 the process of nuclear decay
 elements after #83 are radioactive
 all elements after #92 are synthetic and
decay soon after they are created

42. REVIEW !!!!!!!
 Mass number = # protons + # neutrons
 Atomic number = # of protons
Example
(12 = mass #)
(6= atomic #)
12
6
C

43. REVIEW: ISOTOPES
 most elements have at least one radioactive
isotope
 isotope
 same element with a different number of neutrons
 Example: Carbon-12 (stable)
Carbon-14 (unstable)

44. RADIOACTIVE DECAY
Transmutation
 the process of changing one element into another through
nuclear decay
Radioactive decay
 occurs until a stable nucleus is formed

45. ALPHA DECAY
alpha decay (α)
 releases alpha particle (a helium nucleus)
 a helium nucleus consists of 2 protons and 2 neutrons
 atomic mass is 4
 atomic number is 2
4
2He
 Ex:
238
92
U
234
90
Th +
4
2He

46. BETA DECAY
Beta Decay (β)
 release beta particle it occurs when a neutron breaks down
into 1 electron and 1 proton
the result is an atom with 1 more proton
Ex:
14
6
C
14
7
N+
0 e
-1

47. PRACTICE: ALPHA & BETA DECAY
214
84
222
86
214
82
234
92
Po
Rn
Pb
U
210
82
Pb ______
______
4
2
______
0
1
234
93
Np
He
e
______

48. NUCLEAR REACTIONS
 FISSION
 process of splitting a nucleus into several smaller nuclei

49. NUCLEAR REACTIONS
FUSION:
 Two nuclei with low masses are combined to form one
nucleus of larger mass

50. HALF LIFE
half-life
 the amount of time it takes for
half the nuclei in a sample of the
isotope to decay
the nucleus left after the
isotope decays is called the
daughter nucleus
some half-lives are seconds,
others are millions of years

51. EXAMPLE: HALF LIFE
 Assume a 20g sample of Ba-139 has a
half-life of 86 minutes.
how much Ba-139 remains after 86 minutes?
 after 172 minutes?
 how many ½ lives leave 1.25g of Ba-139?
