Transcript Chemical Reactions Chapter 10 Part I: Counting Atoms How Many Atoms in a Molecule?

Chemical Reactions
Chapter 10
Part I: Counting Atoms
How Many Atoms in a Molecule?
Counting Atoms


Most substances that we encounter are compounds, not
elements.
• A chemical compound is a pure substance formed
from the combination of two or more different elements.


The properties of the compound may be completely unlike those
of the elements that form it.
• The formula for a compound lists the symbols of the
individual elements followed by subscripts which indicate
the number of atoms of that element.

(If no subscript is given, it is understood to be “1.”) E.g., NaCl,
H2O, C12H22O11.
Counting Atoms

A molecular formula gives the actual number of
atoms of each element in a molecule of a
compound.




Hydrogen peroxide H2O2
Water H2O
Glucose C6H12O6
A structural formula uses lines to represent
covalent bonds, and shows how the atoms in a
molecule are joined together:



H—O—O—H
H—O—H
O=C=O
Counting Atoms

Example: How Many Atoms?
C6H12O6
6 C’s + 12 H’s + 6 O’s = 24 atoms
K3PO4
3 K’s+1 P’s + 4 O’s = 8 atoms
C2H5OH
2 C’s+ 5 H’s + 1 O’s + 1 H = 9 atoms
H2O2
2 H’s and 2 O’s = 4 atoms
Counting Atoms with Polyatomic
Ions - Al2(SO4)3
O
O
S
Al
O
O
O
O
S
Al
O
O
O
O
S
O
O
17
=
ATOMS
Counting Atoms with Polyatomic
Ions

When counting atoms with polyatomic ions;

Count number of atoms in one polyatomic ion
o


Ions inside the parentheses
Multiply by number of polyatomic groups in the
molecule (number outside the parenthesis)
Examples:


Al2(SO4)3 - 2 Al’s + 3(1 S + 4 O’s) = 2 + 3(5) = 2+15 =
17 atoms
Mg(NO3)2 – 1 Mg + 2(1 N + 3 O’s) = 1 + 2(4) = 1+8 =
9 atoms
Hydrates

Hydrates are ionic compounds which also
contain a specific number of water molecules
associated with each formula unit. The water
molecules are called waters of hydration.



The formula for the ionic compound is followed by a
raised dot and #H2O
Example: MgSO4•7H2O.
They are named as ionic compounds, followed
by a counting prefix and the word “hydrate”



CuSO4•5H2O copper(II) sulfate pentahydrate
BaCl2•6H2O barium chloride hexahydrate
MgSO4•7H2O magnesium sulfate heptahydrate
(Epsom salts)
How Many Atoms in a Hydrate?
When counting atoms in the hydrate,
count the water atoms also.
 Example:

 CuSO4•5H2O
o
o
o
1 Cu + 1 S + 4 O’s + 5(2 H’s + 1 O’s)
=1+1+4+5(3)
= 6+15 = 21 atoms
 BaCl2•6H2O
o
o
o
1 Ba + 2 Cl + 6(2 H’s + 1 O’s)
= 1+2+6(3)
=3 + 18 = 21 Atoms
Part II: Conservation of Mass
Conservation of Mass

In a normal chemical reaction, the mass of
substances in a closed system will remain
constant, no matter what processes are acting
inside the system.





How ever many atoms a reaction starts with, ends
with the same number.
Atoms don’t change their identity in a chemical
reaction
Number of atoms for EACH ELEMENT STAYS THE
SAME in a chemical reaction
The elements just rearrange their organization
The beginning MASS of the reaction EQUALS the
ending MASS of the reaction
Conservation of Mass

Total Mass stays the same in a chemical
reaction
2g H2 + 16g O2 yields 18g H2O
 Number and Identity of Atoms stays the
same in a chemical reaction
2
H
H
H2 + 1 O2 yields 2 H2O
H
O
H
O
O
H
O
H
H
H
Part III: Writing Reactions
How Do You Write a Chemical
Reaction?
III. Chemical Reactions

Definition – process by which the atoms of
one or more substances are rearranged
 KEY:
new substances are formed
 KEY: No Atoms are Gained or Lost

A chemical reaction is the process by
which atoms of one or more substances
are rearranged into new substances
 Chemical
change occurs
 How do you know?
III. Evidences of a Chemical
Reaction
1)
2)
3)
4)
5)
gas production
light production
temperature change (endo/exothermic)
precipitate formed (solid from 2 liquids)
permanent color change
III. Energy Changes

Energy is stored in compounds as
chemical potential energy
 due

to specific arrangements of atoms.
A chemical reaction changes the potential
energy present.
Energy Changes
•When energy is lost as heat, it is called an
exothermic
reaction
__________________.
These reactions get hotter.
•When energy is gained; heat is added for
a reaction to occur. These are called
endothermic
reactions
______________________,
These reactions get colder.
•Energy in a reaction is shown with:
•ΔH (heat) •Heat
•kJ
•energy
•Joules
III. Chemical Reactions

Representing Chemical Reactions:
– the ‘stuff’ you start with
 An ‘arrow’ which means ‘yields’, or ‘becomes’
 Products – the ‘stuff’ you end up with
 Reactants

Principle of “Conservation of Mass”
applies to chemical reactions.
 Why?
III. Chemical Reactions

Word Equations:
 Reactant-A +
Reactant-B yields Product-AB
 Example:
o
o
Sodium(s) + Chlorine(g) → Sodium Chloride(s)
The small letters in paretheses () indicate the state
of the reactant or product (solid, liquid, gas, or
aqueous solution)




(s) = solid
(l) = liquid
(g) = gas
(aq) = aqueous = dissolved in water
Part IV: Balancing Equations
Applying Conservation of Mass to
Equations
VI. Chemical Equations
Step 1: Write a Skeleton Equation
 Skeleton Equation uses chemical formulas
and symbols instead of words:

 Words:
Sodium + Chlorine gas yields Sodium
Chloride
 Symbols: Na(s) + Cl2(g) → NaCl
 Skeleton Equations are not complete
equations, but are the first step in writing a
complete equation
IV. Chemical Equations

Chemical Equation is BALANCED
 Balanced
means that “conservation of mass”
is upheld
 All atoms in reactants are also in products
o
o
No more, no less
Just rearranged
IV. Chemical Equations

Balancing Equations

Use a number before the compound/element symbol
to indicate how many of them are needed
o
o
o


Called a COEFFICIENT
Written in front of the atom/compound
KEY: Coefficient is a MULTIPLIER
Number of atoms per molecule is SUBSCRIPT
Change ONLY the COEFFICIENTS to balance the
equation
IV. Chemical Reactions

Steps to Balance Equations
1.
2.
3.
4.
5.
6.

Write the skeleton equation
Count the atoms of EACH element in the
reactants
Count the atoms of EACH element in the
products
Change the coefficients to make the number of
atoms of each element equal on both sides of the
equation
Write the coefficients in the lowest possible ratio
Check your work
NEVER CHANGE A SUBSCRIPT
IV. Chemical Equations
1. Write the skeleton equation:
Al + O2 → Al2O3
2. Count Number of atoms for each element on both sides
This is not balanced because the numbers don’t match
3. Multiply coefficients until they match – multiply the entire units
2 Al + O2 → Al2O3
Go to 6 Oxygens
IV. Balancing Equations
2
Al + O2
4
Al +
3
O2
Al2O3
Multiply each atom by 2
2
Al2O3
IV. Balancing Equations 2
The work of balancing a chemical equation is in many ways a series of trials
and errors.
Consider the equation given below.
Does this represent a balanced chemical equation?
N2
+
H2

NH3
IV. Balancing Equations 3
To balance this reaction, it is best to choose one kind of
atom to balance initially. Let's choose nitrogen in this case.
2 Nitrogen Atoms in Reactants requires 2 Ammonia
molecules in Product to balance the nitrogen
N2
+
H2
2NH3
•
•
•
IV. Balancing Equations 2
Once we know what the molecules are (N2, H2, and NH3 in this
case) we cannot change them (only how many of them there
are).
The nitrogen atoms are now balanced, but there are 6 atoms of
hydrogen on the product side
• only 2 of them on the reactant side.
The next step requires multiplying the number of reactant
hydrogen molecules by three to give:
N2
+
H2
3H
2NH3
IV. Don’t Forget: Diatomic Elements

Definition – 7 elements that NEVER occur
as singular atoms (always paired with an
the same or different element)
H2 O2 F2 Br2 I2 N2 Cl2
Ex:
2 HCl + 2K
 2 KCl + H2
IV. Balancing Equations 3
1. Start with an unbalanced equation
2. Draw boxes around the compounds so you don’t mess with the groups
Don’t be threatened by how complex it looks!
IV. Balancing Equations 2
3. Make an element inventory – count number of atoms for
each element on each side of the equation
IV. Balancing Equations 3
4. Write coefficients in front of each of the boxes until the inventory
for each element is the same both before and after the reaction
•Save Oxygen and Hydrogen for last, Treat Polyatomic like an atom.
•Let’s start with Sodium
•We have 2 in products, so I need 2 in reactants
Multiply reactant with sodium by 2 and recount atoms
Element
Reactant
Product
Balanced?
Na
112
2
N
Y
O
112
1
N
Y
H
334
2
N
1
1
Y
SO4
IV. Balancing Equations 3
•Inventory Shows:
•Reactant side has FOUR hydrogen atoms
•Product side has TWO hydrogen atoms
•Using your amazing powers of mathematics
•two hydrogen multiplied two becomes four hydrogen
Element
Reactant
Product
Balanced?
Na
2
2
Y
O
2
112
Y
N
H
4
224
N
Y
SO4
1
1
Y
Helpful Hints

Balance hydrogen and oxygen last

Balance polyatomic ions as a group if present on
both reactants and products


You can consider a polyatomic ion as a single
element
If the balancing starts to get very complex:



Stop
Start over
Select a different atom to balance first.
Example Using PolyAtomics
 Before
MgCl2 + NaOH  Mg(OH)2 + NaCl
1 Mg
2 Cl
1 Na
1 OH
1 Mg
1 Cl
1 Na
2 OH
 After
MgCl2 + 2 NaOH  Mg(OH)2 + 2 NaCl
1 Mg
2 Cl
2 Na
2 (OH)
1 Mg
2 Cl
2 Na
2 (OH)
Types of Chemical
Reactions
Part V
Classifying Chemical Reactions
Synthesis
 Decomposition
 Single replacement
 Double Replacement
 Combustion

Synthesis
– two or more substances
react to form ONE product
 Definition
 Product
is usually bigger or more
complex than either reactant
A + B  AB
Hey baby let’s get jiggy.
Synthesis
 reaction
of two elements
2 Al + ___
3 Cl2  ___
2 AlCl3
___
Al
3+
1Cl
Decomposition
– one substance breaks
down into two or more simpler
products
 definition
AB  A + B
Break yoself fool!
Decomposition

Example reaction:
2 NaN (s)  ___
2 Na (s) + ___
3 N (g)
__
3
2
2 CaO (s)  ___
2 Ca (s) + ___
1 O (g)
__
2
Single Replacement Reactions
– one element replaces
another element in a compound to
form new compound
 Definition
A + BX  AX + B
I’m gon’ dance with yo’ lady
Double Replacement

Defn – exchange of cations between two
ionic compounds
A B + C D  AD + CB
switch
3 possible products of double
replacement reactions
 Precipitate
 Gas
 Water
Reactivity Series (or Activity Series)

More active will replace less active
 Less

active will NOT replace more active
metals
Li K Ca Na Mg Al Mn Zn Fe Ni Sn Pb Cu Ag Au
most active
halogens
least active
F Cl Br I
most active
least active
examples
 aluminum
+ iron (III) oxide
Fe3+ O2-
Stronger?
2 Al
+ 1 Fe2O3
Al
3+
2O
2 Fe + 1 Al O
2 3
examples
 silver
+ copper (I) nitrate Cu1+ NO31Stronger?
Ag +
CuNO3
NO RXN
examples
 fluorine
gas + sodium bromide
Stronger?
1 F
2
+ 2 NaBr
Na1+ F1-
2 NaF + 1 Br
2
examples
 chlorine
gas + hydrogen fluoride
Stronger?
Cl2 +
HF
NO RXN
Example Problem
Li1+

I1-
Ag1+ NO31-
lithium iodide and aqueous silver nitrate
react
Li I + AgNO3
Ag I (s) + LiNO3
Combustion
– compound reacts with O2
 Hydrocarbon – compound w/ only
carbon and hydrogen
 definition
Combustion
 Combustion
of hydrocarbons
ALWAYS produces CO2 and H2O
CxHy + O2
CO2 + H2O
Ex problem

show combustion of propane (C3H8) gas
1 C3H8 + 5 O2
3 CO2 + 4 H2O
I sell propane
and propane
accessories!