Transcript Chapter 2: Formulas and Equations

 Chemical Symbols
 Used to abbreviate the names of elements
 Always start with a capital letter
 Any letters following the first letter MUST be
lowercase
 Chemical symbols with three letters do not yet have
a valid IUPAC name. Their names are functions of
their atomic numbers.
 Some symbols are very easy to remember! 
 Examples
 Carbon = C
 Oxygen = O
 Fluorine = F
 Krypton = Kr
 Argon = Ar
 Some chemical symbols are not easy to remember! 
 Examples:
 Gold – Au (based on the latin word for gold – aurum)
 Sodium – Na (Latin = natrium)
 Mercury = Hg (based on the Greek word
hydrargyrum, which means liquid metal)
 Potassium = K (latin = kalium)
 IUPAC = International Union of Pure and Applied
Chemistry
 A group of scientists from around the world who meet
to discuss the naming of new elements and
compounds.
 This is important since all scientists in the world do
not speak the same language!
 Two identical elements that are chemically combined
 There are SEVEN of them.
 They are:
 Hydrogen – I2
 Nitrogen – N2
 Oxygen – O2
 Fluorine – F2
 Chlorine – Cl2
 Bromine – Br2
 Iodine – I2
 Trick! They make a 7 on the periodic table.
 Start at element 7 (Nitrogen) and make a 7 with your
finger!
 Don’t forget hydrogen! It’s the loneliest element. 
 Although it is not a part of the 7, it IS a diatomic
molecule!
 It is considered to be in its own group on the periodic
table (this will be discussed in more detail later).
 Use chemical symbols and numbers to show both
qualitative and quantitative information about a
substance.
 There are many types, including empirical and
molecular.
 You will be using them throughout this course.
 Example chemical formula: Sulfuric Acid
H2SO4
 Information that can be counted or measured.
 Eg. Mass, height, length, distance.
 Back to our example:
H2SO4
The quantitative information is:
There are TWO Hydrogen atoms
There is ONE Sulfur atom
There are FOUR Oxygen atoms
 Information that cannot be counted or measured.
 E.g., color, ability at a sport, being generally awesome
 Back to our example
H2SO4
Some qualitative information is:
The elements present are Hydrogen, Sulfur and
Oxygen
Sulfuric acid is a liquid at room temperature.
Sulfuric acid is a colorless liquid.
 Subscripts are the tiny numbers that follow a
substance in a chemical formula.
 E.g., CaSO4
 The four is for the oxygen ONLY.
 Subscripts may sometimes follow a special type of
substance in a compound – a polyatomic ion
 When this happens, parentheses must surround the ion.
 E.g., Ca3(PO4)2
 The first 3 is for the calcium. The four is for the oxygen.
However, the 2 outside of the parentheses is for both the
phosphorus AND the oxygen!
 Coefficients are large numbers, written before a
chemical formula, which denotes the number of moles
of that substance.
 E.g., 10 H2O
There are 20 atoms of hydrogen, and 10 atoms of
oxygen!
List the qualitative and quantitative information for the
following substances:
1) NH3
2) (NH4)3PO4
3) 4 Au
4) 7 C6H12
5) 8 Mg(NO3)2
 There are two general types of chemical formulas:
 Empirical Formulas – chemical formulas in which the
ratio of all atoms in the compound is in its simplest
form.
 Molecular Formulas – chemical formulas in which the
ratio of all atoms in the compound is not in its simplest
form.
 Trick for later – all molecular formulas represent covalently
bonded substances (don’t worry about this until later)
H2O
Ratio of hydrogen atoms to oxygen atoms is 2:1
NaCl
Ratio of sodium atoms to chlorine atoms is 1:1.
CaSO4
Ratio of calcium atoms to sulfur atoms to oxygen
atoms is 1:1:4
Butane – C4H10
If you divide both subscripts by 2, you get C2H5.
C2H5 is the empirical formula for C4H10.
Hydrogen peroxide – H2O2
If you divide both subscripts by 2, you get HO.
HO is the empirical formula for H2O2.
Glucose – C6H12O6
If you divide both subscripts by 6, you get CH2O
CH2O is the empirical formula for C6H12O6.
 Ions, simply, are charged particles that result from an
UNEQUAL number of protons and electrons.
 Why does this happen!? BONDING – it’s OWN chapter!
 You represent the charge of at atom via a superscript,
which is a number written like a “raise to the power of”
in math.
 E.g., calcium ion – Ca2+
 Oxygen ion – O2 Aluminum ion – Al3+
 Polyatomic Ions are ions that consists of more than
one atom.
 All of the ones you are responsible for can be found on
TABLE E in your reference tables.
 The vast majority of polyatomic ions and in –ATE or –
ITE.
 Get used to this table…you’ll be using it ALL year!
 Examples of Polyatomic Ions
 Sulfate – SO42 Phosphate – PO43 Nitrite – NO2 Ammonium – NH4+ (note – no ate or ite ending)
 Hydroxide – OH- (note no ate or ite ending)
 Peroxide – O22-
 Hydrates are ionically bonded substances, in crystal
form, in which molecules of water are trapped inside
of a crystalline lattice.
 Think of it like little tiny water molecules being caught
in a net.
 Example of a hydrate: copper (II) sulfate or cupric
sulfate:
CuSO4
∙ 5H O
2
 The dot in the middle does not mean multiplication!
 The dot is used as notation to separate the anhydrous
crystal from the water.
 Depending on the hydrate, they may be pretty colors:
 Gets ugly.
 How many atoms of each element in
 10 CuSO4
∙ 5H O
Copper – 10
Sulfur – 10
Oxygen – 45!
Hydrogen - 10
2
 CuCl2 • 2H2O
 Naming compounds ranges from being really easy to
slightly difficult, based on the type of chemical
compound.
 We’ll start off easy – with naming compounds.
 At this point, a binary ionic compound consists of a
metal and nonmetal.
 Metals are found to the left of the bolded stairs on the
Periodic Table of the Elements.
 Nonmetals are found to the right of the bolded stairs
on the Periodic Table of the Elements.
 Binary ionic compounds consist of only two elements!
 NAME = first element name + second element name
ending in –ide
 The first element in all compounds is considered to be
“positive” and the second element is considered to be
“negative.”
 LiF = lithium fluoride
 CaO = calcium oxide
 AlBr3 = aluminum bromide
 Rb2S = rubidium sulfide
 Nonbinary ionic compounds have more than two
elements in the compound.
 If this is the case, the compounds MUST have a
polyatomic ion from Table E present!
 NAME = first substance name + second substance
name
 NOTE: If the second substance is an element, the
ending must still change to –ide!
 LiNO3 = lithium nitrate
 CaCO3 = calcium carbonate
 Fr2SO4 = francium sulfate
 Ba3(PO4)2 = barium phosphate
 Binary covalent compounds consist of two nonmetals
that are chemically combined.
 You must use the Prefix Table (Table 2-5 on p26 of your
workbook) to figure out how many atoms of each
element are in the cmpound.
 The names for binary covalent compounds still end in
–ide.
Memorize this.
Memorize this.
Seriously.
 P4H10 = tetraphosphorus decahydride
 SO3 = sulfur trioxide
 O2F2 = dioxygen difluoride
 CO = carbon monoxide
 CO2 = carbon dioxide
 Depending on the type of compound (ionic or
covalent), there are different rules to follow.
 We’ll start easy again…
 Use the prefix table to write the formula for the
compound!
 Write the formula for carbon tetrabromide.
 Answer: CBr6
 Write the formula for dichlorine monoxide.
 Answer: Cl2O
Memorize this.
Seriously.
For real!
 Remember: a binary ionic compound consists of one
metal and one nonmetal.
 You need to start utilizing oxidation numbers from the
Periodic Table of the Elements!
Example Ionic Compound:
Aluminum iodide
Question: What is the chemical formula for this
compound?
 Write out the symbols of the elements right next to
each other.
1) AlI
 Look up the oxidation numbers for each element in
the compound, and write them above each element.
Al+3I-1
 CRISS CROSS!
 Criss cross the oxidation numbers in the compound.
They now become subscripts. Drop any signs
associated with the numbers.
Al+3I-1
AlI3
 All ionic compounds must be written in empirical
form!
 So, if the oxidation states for the two elements are
numerically equivalent, they will cancel each other out
– DO NOT CRISS CROSS!
 Eg: strontium sulfide
+2
-2
Sr S
= SrS
1)
Calcium oxide
2) Rubidium bromide
3)
Potassium sulfide
4) Magnesium oxide
 A nonbinary ionic compound has a polyatomic ion
present in it.
 Written in the same fashion as binary ionic
compounds, but polyatomic ions have CHARGES (not
oxidation states).
 These are the numbers you criss-cross!
Write the chemical formula for nickel (II) sulfate.
Nickel = Ni
Sulfate = SO4
1) Write everything out: NiSO4
Write the chemical formula for nickel (II) sulfate.
Nickel = Ni
Sulfate = SO4
2) Look up oxidation numbers / charges:
Ni+2SO4-2
Remember: Nickel is +2 because of the STOCK SYSTEM!
Write the chemical formula for nickel (II) sulfate.
Nickel = Ni
Sulfate = SO4
3) Criss Cross!
Ni+1SO4-2
Write the chemical formula for nickel (II) sulfate.
Nickel = Ni
Sulfate = SO4
Answer!
NiSO4
(the charges cancel out)
Write the chemical formula for aluminum nitrate.
1) Write everything out: AlNO3
Write the chemical formula for aluminum nitrate.
2) Look up oxidation numbers / charges
Al+3NO3-1
Write the chemical formula for aluminum nitrate.
2) Criss Criss!
Al+3NO3-1
Write the chemical formula for aluminum nitrate.
Answer!
AlNO33
WHAT!? Thirty Three oxygen atoms!? That doesn’t
make sense! AAAAARGH!
Write the chemical formula for aluminum nitrate.
 When you criss-cross, and have more than one
polyatomic ion, you MUST put that ion in parentheses
to separate the subscripts.
Write the chemical formula for aluminum nitrate.
Actual Answer!
Al(NO3)3
Now there are a total of NINE oxygen atoms!
1)
Aluminum phosphate
2) Ammonium sulfate
3)
Lithium carbonate
4) Sodium hydroxide
 Chemical Equations are written to represent both
chemical and physical changes.
 They are the combination of chemical formulas
arranged to convey an idea.
 New symbols are also used which you need to learn!
 A physical change is a change in which a brand




spankin’ new substance does NOT result after the
change.
You still have the original substance, but it’s just in a
new form.
The substance retains its original properties.
Can be represented by chemical equations!
Physical changes are typically PHASE changes.
 PHysical = PHase
 Phase Changes: melting, freezing, evaporation, etc.
 Ice is made of water. Liquid water is also made
of…water! When ice melts, it’s still water after the
change, just in a different form.
 Painting: Your house is still your house if you repaint
it.
 Example: Melting of Ice
H2O(s) + energy  H2O(l)
 A chemical change is a change in which a brand new
substance forms as a result of the change.
 The new substance(s) formed attain NEW properties!
 Can be represented by chemical equations.
 Key words are typically RUSTING, BURNING,
COMBUSTING, REACTING
 Rusting:
 Combustion:
2H2(g) + O2(g)  2H2O(g) + energy
2H2(g) + O2(g)  2H2O(g) + energy
 Substances to the left of the arrow are reactants. These
are the substances that interact to cause a chemical
reaction to occur.
2H2(g) + O2(g)  2H2O(g) + energy
 Substances to the right of the arrow are products.
These are the substances that result from the chemical
reaction between the reactants.
2H2(g) + O2(g)  2H2O(g) + energy
 Plus signs are treated as commas. The reactants and
products are merely lists of substances. The order on
each side of the arrow does not matter.
 The arrow in the reaction separates each side of the
reaction. It is read as “yields,” or “makes.”
2H2(g) + O2(g)  2H2O(g) + energy
 Letters in parentheses represent a physical property of
the substance preceding it.
 (s) = solid
 (l) = liquid
 (g) = gas
 (aq) = aqueous – The preceding substance is dissolved in
water! This means it is a HOMOGENEOUS MIXTURE!
2H2(g) + O2(g)  2H2O(g) + energy
 All chemical reactions deal with energy.
 If energy is a product (right side of the arrow), the
reaction is considered to be exothermic.
 Energy EXits the reaction.
 If energy is a reactant (left side of the arrow), the
reaction is considered to be endothermic.
 Energy ENters the reaction.
2H2(g) + O2(g)  2H2O(g) + energy
 This is an exothermic reaction. Energy is released
after the reaction takes place.
H2O(s) + energy  H2O(l)
 This reaction is endothermic. Energy is absorbed as
the reaction takes place.
 In order to break a bond, energy MUST be absorbed to
overcome the attractive forces between atoms
competing for electrons.
 BREAKING A BOND IS ALWAYS ENDOTHERMIC.
 Alternately, when a bond is formed, energy is released
into the environment.
 FORMING A BOND IS ALWAYS EXOTHERMIC.