Download Chemical Formulas and equations

Chemical Formulas and
Equations
Do Now
• What is a mole in chemistry?
• What was Avogadro's number? What did it
represent?
• What is molar mass?
Do Now
• What is a mole in chemistry?
• Remember: Mole (mol) = the SI unit for
amount
• What was Avogadro's number? What did it
represent?
• Remember: Avogadro's number = the
number of particles in a mole = 6.02 x 1023
• What is molar mass?
• Remember: Molar mass = mass in grams of
one mole of an element or compound
Remember
• You can convert from moles to particles,
or moles to mass, and vice versa.
• Ex: How many particles do you have if you
have 2.5 mol of sulfur?
• Ex: How many grams of carbon do you
have if you have 2.44 x 1022 atoms?
* You must use conversion factors to convert from number of
atoms to moles and then to grams.
Formula Mass
• The molar mass is numerically equal to
the atomic mass of monatomic elements
and the formula mass of compounds and
diatomic elements.
• Ex: Find the formula mass of KBr
• Ex: Find the formula mass of H2O
Chemical Formula
• Chemical formulas indicate the relative
number of atoms of each kind in a
chemical compound.
• Identify the number of atoms in each of the
following compounds:
1. KCl= _____________________
2. C6H12O6=__________________
3. NH3=__________________
Chemical Formulas
• Formulas for covalent compounds show
both the elements and the number of
atoms of each element in a molecule.
• Formulas for ionic compounds do not
show numbers of atoms, but show the
simplest ratio of cations and anions.
Chemical Formulas
• The meaning of formulas do not change
when polyatomic ions are involved.
• Polyatomic ions = a group of covalently
bond elements that behave as a single ion
Finding the number of atoms in a
compound with polyatomic ions
Compound
1. MgCO3
2. Al(ClO4)3
What does it look like
expanded?
Numbers of
atoms
Compound
3. Zn3(PO4)2
4. (NH4)2S
What does it look like
expanded?
Numbers of
atoms
Molar Mass
• Formulas can be used to calculate molar
mass
• Ex: ZnCl2
• Ex: ZnSO4
Do Now
• Do you see a pattern in the following
formulas?
NH4NO2
NH2O
Law of Definite Proportions
• Law of Definite Proportions states that
every pure substance always contains the
same elements combined in the same
proportions by weight.
• For example: H2O, will always have the
same percent by weight (11.2% H and
88.8% O)
Percentage Composition
• Percentage composition = the percentage
by mass of each element in a compound
• Percentage composition helps verify a
substance’s identity and can be used to
compare the ratio of masses contributed
by the elements in two substances
• Ex: Fe2O3
FeO
69.9% Fe
77.7% Fe
30/1% O
22.3% O
Review Formulas
• Molecular Formula = represents the
number and kind of atoms in a molecule
(not necessarily the smallest whole
number ratio).
• Structural Formula = indicates twodimensional arrangement of the bonds
and lone pairs of electrons in a molecule.
Empirical Formulas
• Empirical formula = a chemical formula
that shows the composition of a compound
in terms of the relative numbers and kinds
of atoms in the simplest ratio.
Empirical Formula for Molecules
(have covalent bonds)
• The chemical formula that represents the simplest (lowest) atomic
ratio in which elements can combine.
• Formulas for molecular compounds are NOT NECESSARILY the
empirical formula
Molecular Formula
Empirical Formula
C2H6
C6H12O6
C4H8
C3H8
C6H10
C4H6
C2H4
C5H10
Empirical Formula for Ionic Compounds
(Always written in empirical form)
• The chemical formula that represents the simplest (lowest) atomic
ratio in which elements can combine.
• Formulas for ionic compounds are the empirical formula because
they are ALWAYS written in the expressed as the lowest possible
ratios.
Empirical Form
Ca2O2
Fe2O2
Pb2O4
Mg2O2
Empirical Formula for Ionic
Compounds with Polyatomic Ions
Empirical Form
Fe2(CO3)2
Fe2(HSO4)2
Zn2(SO4)2
Pb2(SO3)2
Empirical or Molecular
Formula
C6H12O6
LiNO3
H2O2
C11H22O11
If written in molecular
formula what is the
empirical formula?
Type of Bond
Finding Empirical Formulas
• You can find the empirical formula from
percentage composition!
– Step 1: Convert to mass in grams (assume
you have 100g of the given substance)
– Step 2: Convert from grams to moles (using
the molar mass conversion factor)
– Step 3: Reduce the molar ratio to the simplest
whole-number ratio by dividing by the smaller
amount
– Step 4: Round to whole numbers and insert
subscripts
Finding Empirical Formulas
• Ex: A given liquid has 60.0% C, 13.4% H,
and 26.6% O by mass. Calculate the
empirical formula.
– Step 1: Convert to mass in grams (assume
you have 100g of the given substance)
60.0% C x 100g = 60.0g C
13.4% H x 100g = 13.4g H
26.6% O x 100g = 26.6g O
Finding Empirical Formulas
– Step 2: Convert from grams to moles (using
the molar mass conversion factor)
60.0g C x 1mol = 5.00mol C
12.01g C
13.4g H x 1mol = 13.3mol H
1.01g H
26.6g O x 1mol = 1.66mol O
16.00g O
C5H13.3O1.66 ?
Finding Empirical Formulas
– Step 3: Reduce the molar ratio to the simplest
whole-number ratio by dividing by the smaller
amount
C5H13.3O1.66 ?
5.00mol C / 1.66 = 3.01mol C
13.3mol H / 1.66 = 8.01mol H
1.66mol O / 1.66 = 1.00mol O
– Step 4: Round to whole numbers and insert
subscripts
C3H8O
Practice
• What is the empirical formula of a
compound that is 78.6% B and 21.4% H?
Practice
• What is the empirical formula of a
compound containing 32.38% Na, 22.65%
S, and 44.99% O?
Practice
• What is the empirical formula of a
compound containing 26.56% potassium,
35.41% chromium, and the remainder is
oxygen?
Practice
• In a 10.150g sample, 4.433g are
phosphorous and the rest is oxygen. What
is the empirical formula for this
compound?
Practice
• A compound contains 0.606g nitrogen and
1.390g oxygen. What is the empirical
formula of the compound?
Do Now
• What was the difference between the
empirical formula and molecular formula?
Do Now
• What was the difference between the
empirical formula and molecular formula?
• Empirical formula = smallest possible
whole number ratio of elements
• Molecular formula = the actual formula of a
molecule
Molecular Formulas
• For ionic compounds, the molecular
formula is the same as the empirical
formula.
• For molecular compounds, the molecular
formula is a whole number multiple of the
empirical formula.
*Both formulas are just different ways of representing the
composition of the same molecule.
Molar Mass
• In order to determine the molecular
formula, you must know the molecular
mass!
• The molar mass of a compound is equal to
the molar mass of the empirical formula
time a whole number, n.
n(empirical formula) = molecular formula
Molecular Formula Examples
• Formaldehyde, acetic acid, and glucose
each have the same empirical formula,
CH2O
– For formaldehyde, n = 1
– For acetic acid, n = 2
– For glucose, n = 6
Determining Molecular Formula
• Step 1: Find the molar mass of the
empirical formula using the molar masses
of the elements from the periodic table
• Step 2: Solve for n, the factor multiplying
the empirical formula to get the molecular
formula
n = experimental molar mass of compound
molar mass of the empirical formula
• Step 3: Multiply the empirical formula by
this factor to get the molecular formula
Determining Molecular Formula
• The empirical formula of a compound is
P2O5. The experimental molar mass is 284
g/mol. Determine the molecular formula.
– Step 1: Find the molar mass of the empirical
formula using the molar masses of the
elements from the periodic table
2 x molar mass of P = 2(30.97) = 61.94g/mol
5 x molar mass of O = 5(16.00) = 80.00g/mol
molar mass of P2O5 = 141.94g/mol
Determining Molecular Formula
• The empirical formula of a compound is
P2O5. The experimental molar mass is 284
g/mol. Determine the molecular formula.
– Step 2: Solve for n, the factor multiplying the
empirical formula to get the molecular formula
n = experimental molar mass
molar mass of empirical formula
n = 284 g/mol
141.94 g/mol
n=2
Determining Molecular Formula
• The empirical formula of a compound is
P2O5. The experimental molar mass is 284
g/mol. Determine the molecular formula.
– Step 3: Multiply the empirical formula by this
factor to get the molecular formula
n(empirical formula) = 2(P2O5) = P4O10
Determining Molecular Formula
• You can verify your answer by finding the
molar mass of the molecular formula and
compare it to the experimental molar mass
4 x molar mass of P = 4(30.97g/mol) = 123.88g/mol
10 x molar mass of O = 10(16.00g/mol) = 160.0g/mol
molar mass of P4O10 = 283.88g/mol
Practice
• What is the molecular formula of a
compound with the empirical formula BH3
and molecular mass of 28g?
Practice
• What is the molecular formula of a
compound with a molecular mass of 34g
that consists of 0.44g H and 6.92g O?
Practice
• What is the empirical formula of a
compound that contains 65.5% carbon,
5.5% hydrogen, and 29.0% oxygen? What
is the molecular formula if the molecular
mass is 110g?
Do Now
• What does percentage mean?
• What does a percentage represent?
Review
• Chemical formulas allow scientists to
calculate a number of characteristics
values for a given compound.
• Chemical formulas represent the number
and kind of atoms in a molecule.
• If you know the chemical formula, then you
can calculate the percentage composition.
Review
• Formula mass = the sum of the average
atomic masses of all the atoms
represented in the formula
• Formula mass is numerically equal to the
molar mass or gram formula mass
• Ex: Find the formula mass of KClO3
• Ex: What is the molar mass of Ba(NO3)2?
Percent Composition
• Percentage composition = the percentage
by mass of each element in a compound
Finding Percent Composition
• From the subscripts, you can determine
the mass contributed by each element and
add these to get the molar mass.
• Divide the mass of each element by the
molar mass.
• Multiply by 100 to find the percentage
composition of that element
Finding Percent Composition
• CO2
1mol x 12.01 g/mol = 12.01 g C
+2mol x 16.00 g/mol = 32.00 g O
mass of 1mol CO2 = 44.01 g
Finding Percent Composition
% C = 12.01 g C x 100 = 27.29%
44.01 g CO2
% O = 32.00 g O x 100 = 72.71%
44.01 g CO2
Finding Percent Composition
• CO
1mol x 12.01 g/mol = 12.01 g C
+1mol x 16.00 g/mol = 16.00 g O
mass of 1mol CO = 28.01 g
Finding Percent Composition
% C = 12.01 g C x 100 = 42.88%
28.01 g CO
% O = 16.00 g O x 100 = 57.71%
28.01 g CO
Percent Composition
• Percentage composition helps verify a
substance’s identity and can be used to
compare the ratio of masses contributed
by the elements in two substances
• Ex: Fe2O3
FeO
69.9% Fe
77.7% Fe
30/1% O
22.3% O
Practice
• Calculate the percentage composition of
Cu2S, a copper ore called chalcocite.
Practice
• Calculate the percent of both elements in
sulfur dioxide
Practice
• Calculate the percentage composition of
ammonium nitrate, NH4NO3
Hydrates
• Hydrates = salts that have crystallized
from water solution
• In the process of crystallization, the water
molecules bind to the salt to form
hydrates.
• Hydrates are represented as follows:
Na2CO3 ● 10H2O
(sodium carbonate decahydrate)
Practice
• Calculate the molar mass of
Na2CO3 ● 10H2O
Practice
• Calculate the percent composition of water
in Na2CO3 ● 10H2O
Do Now
• Take out Reference Tables and open to
Table E – QUIZ ON POLYATOMIC IONS
ON WEDNESDAY
• Turn to Table S – Where is the oxidation
number noted?
• How many aluminum atoms combine with
how many sulfur atoms?
How do we figure this out?
OXIDATION NUMBER
• Oxidation number = the charge an atom
would acquire if all its bonds were treated
as ionic bonds.
• To determine how many atoms combine
with one another in a compound we must
determine each element’s OXIDATION
NUMBER.
Predicting Ionic Charges
Group 1: Lose 1 electron to form 1+ ions
H+
Li+
Na+
K+
Predicting Ionic Charges
Group 2: Loses 2 electrons to form
2+ ions
Be2+
Mg2+
Ca2+
Sr2+
Ba2+
Predicting Ionic Charges
B3+
Al3+
Ga3+
Group 13: Loses 3
electrons to form
3+ ions
Predicting Ionic Charges
Group 14:
Lose of 4
electrons or gain
of 4 electrons?
Neither! Group 14
elements rarely form
ions.
Predicting Ionic Charges
N3- Nitride
P3- Phosphide
As3- Arsenide
Group 15: Gains 3
electrons to form
3- ions
Predicting Ionic Charges
O2- Oxide
S2- Sulfide
Se2- Selenide
Group 16:
Gains 2
electrons to form
2- ions
Predicting Ionic Charges
F1- Fluoride
Br1- Bromide
Cl1-Chloride
I1- Iodide
Group 17:
Gains 1
electron to form
1- ions
Predicting Ionic Charges
Group 18:
Stable Noble gases
do not form ions!
Predicting Ionic Charges
Groups 3 - 11: Many transition elements
have multiple oxidation states.
Iron(II) = Fe2+
Iron(III) = Fe3+
Predicting Ionic Charges
Groups 3 - 11: Some metals
have only one possible oxidation state.
Silver = Ag+
Zinc = Zn2+
Rules for assigning Oxidation
Numbers
1. The atoms in a pure element have an oxidation number of zero.
2. Alkali metals always have an oxidation number of +1; alkaline
earth metals always have an oxidation number of +2.
3. Fluorine always has an oxidation number of -1.
4. Oxygen has an oxidation number of -2 in almost all compounds.
Exceptions are in compounds with a halogen, when it has an
oxidation number of +2, and in peroxides (H2O2), when it has
an oxidation number of -1.
5. Hydrogen has an oxidation number of +1 in almost all
compounds except when combined with a metal when it has an
oxidation number of -1.
6. The sum of all the oxidation numbers in a neutral compound is
zero.
7. The sum of all the oxidation numbers in a polyatomic ion is
equal to the charge of the ion.
ASSIGNING OXIDATION
NUMBERS
• A compound has a total charge of ZERO so set
your equation equal to ZERO.
• Assign the variable X to your unknown oxidation
number.
• Given a compound, find the oxidation number of
every element you know for certain. Then solve
for others using algebra.
• KMnO4
• CaCO3
Do you notice a pattern with the
elements in these compounds?
H2+1 O1-2
Mg1+2F2-1
Al2+3S3-2
Cu3+1P1-3
Ca1+2Cl2-1
Fe2+3O3-2
CRISS-CROSS METHOD
to determine the chemical formula
1. Write the symbols for the elements side by side.
2. Write the oxidation states of each element to the
top right of the symbol. When the nonmetal is
combines with a metal, the oxidation state will
always be the first number (the negative one) in the
list of oxidation states.
3. Criss cross the charges DOWN and use the
absolute values (-2 becomes 2).
4. Check to make sure the subscripts are the lowest
ratio.
Practice Criss-Cross Method
1. Na and S
2. K and P
3. Al and S
4. Mg and Br
5. Al and O
What about Polyatomic Ions?
• Write down the only cations on Table E:
(positively charges ions)
• Write down the only polyatomic ions that
end in –ide.
Polyatomic Ions Chart (Table E)
Find the charge of the polyatomic ion using Table E
Put parenthesis around the polyatomic ion.
1. PO4
7. HSO4
2. CO3
8. ClO4
3. SO3
9. CN
4. NH4
10. OH
5. ClO
11. S2O3
6. ClO2
12. SCN
Polyatomic Ions
• If you see a group of atoms together with a charge it
is a polyatomic ion from Table E.
• Put the polyatomic ion within Parenthesis.
• Find the charge of the polyatomic ion.
• Use the criss-cross method to determine subscripts.
Na +1 CO3-2
How many atoms are present in the compound?
NH4 +1 S-2
How many atoms are present in the compound?
Rules for writing the formula of
the compound composed of ions
1. Place all Polyatomic Ions in Parenthesis
(Table E)
2. Determine all oxidation numbers of
elements and polyatomic ions
3. Use Criss-Cross Method
4. Reduce to Empirical Form
Practice
1. NH4 S
2. Na NO3
3. Cu Br
4. Al SO4
5. Fe CO3
6. Pb PO4
7. Ag ClO
8. Ca F
9. NH4 SO3
10. Cu OH
11. Ni I
12. Zn SO4
13. Pb ClO2
14. H I
15. Fe HSO4
16. Cu CO3
17. NH4 O
18. Ag S
19. Al ClO4
Do Now
• What is the name of H2O?
• What is the name of NaCl?
• What is the name of NH4?
Naming Compounds
• Many chemical compounds have common
names.
• H2O = water
• NaCl = table salt
• NH4 = ammonia
Review
• Chemical formula = the relative number of
atoms of each kind in a chemical
compound
• Molecular formula = the number and kind
of atoms in a molecule (not necessarily the
smallest whole number ratio)
NAMING IONIC COMPOUNDS
• An Ionic compound can quickly be
determined if a METAL is bonded to a
NONMETAL.
• Naming Ionic compounds with metals that
have only ONE oxidation state is fairly
simple.
Of the following metals listed below,
check off all the elements that have
more than one oxidation state:
a. Mn___
b. Zn ___ c. K ____
d. Pb___
e. Mg___
f. Au ___ g. Ag ___
h. Ga__
i.
j. Li ____ k. U ____
l. Cr ___
Sn __
Naming Ionic Compounds that
contain metals with 1 oxidation state
• Binary Ionic Compounds = ionic
compounds with only 2 different
elements.
• Name the metal and end the nonmetal in
–ide.
• For example: CaBr2 = _______________
Naming Binary Ionic Compounds
• Naming salts is very easy, because they are binary ionic
compounds (made up of two elements).
– The cation is named by borrowing the name of the
element.
– The anion named by combining the name of the
element with an –ide ending.
• The name of compound is made up of both the cation
and anion name
– Ex: NaCl = sodium chloride
– Ex: ZnS = zinc sulfide
– Ex: K2O = potassium oxide
– Ex: Mg3N2 = magnesium nitride
– Ex: Al2S3 = aluminum sulfide
Name the following Binary
compounds:
1. MgO= ________________________
2. CaCl2 = _______________________
3. AlBr3 = _______________________
4. Ag3N = _______________________
5. Al2O3 = ______________________
6. LiI = _________________________
7. BaF2 = _______________________
8. Zn2C = _______________________
9. Ba3N2 = ______________________
10.CdO = _______________________
11. Ga2S3 = _______________________
12. K3N = ________________________
13. SrO = ________________________
What about Polyatomic ions?
(Table E)
• These are ions consisting of more than
one atom. The names of polyatomic ions
end in –ate or –ite, except for two
(ammonium and cyanide)
Polyatomic ion chart (Table E)
Ternary Ionic Compounds
•
•
•
•
Ionic Compounds with 3 different elements.
They usually contain Polyatomic Ions (Table E)
Name the following Polyatomic Ions:
a. NO3- _________ d. SO42- ___________
• b. ClO2- _________ e. SO32- __________
• c. CO32- _________ f. SCN- ____________
Naming Ternary Compounds
• Name the metal and then name the
polyatomic ion (if it has a negative
oxidation number).
• For example: KNO3 = ________________
Naming Ternary Compounds
• Name the polyatomic ion (if it has a
positive oxidation number) and then name
the nonmetal.
• For example: NH4Cl = ________________
Naming Ternary Compounds
• If there are two polyatomic ions, name the
positive polyatomic ion first and then name
the negative polyatomic ion.
• For example: NH4NO3 = ______________
• DO NOT CHANGE THE ENDINGS OF
POLYATOMIC IONS!
Name the following Ternary
compounds:
1. NaC2H3O2 _____________________
2. AgHCO3 _______________________
3. LiNO2 _________________________
4. Ga2(S2O3)3 ____________________
5. Ca3(PO4)2 _____________________
6. ZnSO3 __________________________
7. KClO3 ___________________________
8. Al(OH)3 _________________________
9. RbSCN __________________________
10. SrCO3 _________________________
Naming Ionic Compounds with Metals with
Multiple oxidation states (multiple charges):
Using the STOCK SYSTEM
1. Determine the oxidation state of the metal
in the compound.
2. Name the metal, put the oxidation state in
ROMAN NUMERALS in parenthesis and
end the nonmetal in –ide.
Review: Find the Formula
Criss-Cross
Pb+4 O-2
Cu+2 (SO4) -2
Sn+2(CO3)
-2
Find Empirical
formula
Work Backwards: Start with the Empirical Formula to determine
the Oxidation state of a Metal with Multiple Oxidation States
Empirical
Formula
Fe1O1
Write in the oxidation
number for the nonmetal
or polyatomic ion you are
sure of and criss-cross.
Fe O
Fe1(SO4)1
Fe (SO4)
Cu1(SO4)1
Cu (SO4)
Non-reduced form with
Oxidation States
Work Backwards: Start with the Empirical Formula to determine
the Oxidation state of a Metal with Multiple Oxidation States
Empirical
Formula
Write in the oxidation
number for the nonmetal
or polyatomic ion you are
sure of and criss-cross.
Sn1(SO3)1
Sn (SO3)
Mn1(SO4)2
Mn (SO4)
Cr1(PO4)2
Cr (PO4)
Non-reduced form with
Oxidation States
Name the following compound
using the Stock System:
1. Fe O ____________________________
2. Fe Cl2 ___________________________
3. Cu SO4 __________________________
4. Pb Cl2 ___________________________
5. Pb O2 ___________________________
6. Cu3(PO4)2 ________________________
7. Cu2 S ___________________________
8. Fe2(CrO4)3 ______________________
19. Sn CO3 _________________________
10. Sn F4 ___________________________
Name each of the following
compounds, use Roman Numerals
only when necessary.
Put a check next to every compound that begins
with a metal with more than 1 oxidation state.
Put parenthesis around all the polyatomic ions.
1. NH4 Cl _________________________
2. Pb SO4 _________________________
3. Co Cl3 __________________________
4. Ba (NO3)2 ______________________
5. Co2 (SO3)3 _____________________
6. KH ____________________________
7. NH4 F _________________________
8. K2Cr2O7 _______________________
9. Cu S __________________________
10. Cu ClO2 ______________________
11. Ag NO3 _______________________
12. Fe Cl3 ________________________
13. Cr F2 _________________________
14. Na Cl _________________________
15. Fe PO4 ______________________
16. Li F _________________________
17. Fe F3 _______________________
18. Al (OH)3 _____________________
19. Mg I2 ________________________
20. Fe Cl3 _______________________
Do Now
• What types of bonds are in molecular
compounds? _____________________
• How can we tell if a formula has a covalent
bond? ___________________________
Review Molecular Compounds
• Molecule = a neutral compound held
together by covalent bonds
• Molecules may consist of identical atoms
bonded together (O2) or different atoms
bonded together (H2O)
Naming Covalent Compounds
• Covalent compounds are named in a
similar way to ionic compounds
– The first element in the formula is usually
written first in the name
– The second element has an –ide ending
• Ex: SO2 = sulfur oxide
– However, this is not completely correct…
Naming Covalent Compounds
• Since multiple covalent compounds can be
made from the same elements, the name
must distinguish them as different.
– Prefixes are used to indicate the number of
atoms of each element in the molecule.
• Ex: SO2 = sulfur dioxide
• Ex: SO3 = sulfur trioxide
Naming Covalent Compounds
• Most nonmetals have more than 1
oxidation state, therefore you can use the
Stock System (Roman Numerals) or the
Prefix System.
• The Prefix System includes:
MonoDiTriTetraPenta-
HexaHectaOctaNonaDeca-
Naming Molecular Compounds
• According to the number of atoms of each
element, state the prefix for the number of
each atom before the name of the element
and end the nonmetal in –ide.
• For example:
P2O5 = ______________
CO2 = ______________
Naming Molecular Compounds
• Name the first element. Use a prefix ONLY
if there is more than one.
• Name the second element. ALWAYS use
a prefix. Change the ending to –ide.
• The prefix mono only needs to be used for
the second half of the compound NOT the
first element.
Name the following covalent
compounds:
1.
CI4 ________________________________
2.
PCl5 ______________________________
3.
SI6 _______________________________
4.
P2S6 ______________________________
5.
N3O4 ______________________________
6. SO2 ____________________________
7. N2O4 ___________________________
8. CO _____________________________
9. NF3 ____________________________
10. ICl5 ___________________________
11. H2S3 __________________________
12. N2O3 _________________________
13. ClF7 __________________________
14. SO3 ___________________________
15. NI5 ___________________________
16. BN2 __________________________
17. P2O5 _________________________
18. IF7 ___________________________
Do Now
• What is the chemical formula for
dihydrogen oxide?
• What is the chemical formula for acetate?
Writing Formulas
• The chemical formula of a compound can
be determined from the chemical name!
General Steps to Determining
Chemical Formulas:
• 1. Determine what elements and/or
polyatomic ions are in the compound.
• 2. Write the symbols for each substance
with oxidation states of each substance to
the top right of the symbol.
– If there is a metal with more than one
oxidation state, it will be indicated as roman
numerals in parenthesis after the element.
– If there is a nonmetal with more than one
oxidation state, use the negative number.
General Steps to Determining
Chemical Formulas (continued):
• 3. All oxidation states in a neutral compound
add up to ZERO, so figure out how many of
each substance you need to make all
oxidation states add up to ZERO.
• 4. Put these numbers as subscripts and write
the chemical formula.
• 5. If you have more than one polyatomic ion,
make sure you put the symbols in
parenthesis.
• 6. Check to make sure the subscripts are the
lowest ratio.
Polyatomic Ions (Table E)
Writing Formulas for Ionic
Compounds
• All ionic compounds must be in empirical form.
(Reduced Form)
• If the compound ends in –ide, most likely it is a
binary compound. Except for cyanide and
hydroxide.
• If the substance ends in –ate or –ite it contains a
polyatomic ion. Put the polyatomic ion in
parenthesis with oxidation numbers indicated
before you criss-cross. Put the oxidation states
on top and criss-cross the numbers.
•
For example: Aluminum Sulfide =
Write the formulas for each of the
following compounds:
(Make sure you reduce it to empirical form)
1. Aluminum Chloride ___________________
2. Silver Phosphate _____________________
3. Lithium Hydride ______________________
4. Magnesium Acetate __________________
5. Potassium Sulfite _____________________
6. Zinc Thiosulfate ____________________
7. Strontium Nitride ____________________
8. Calcium Oxide _______________________
9. Gallium Oxalate ______________________
10. Ammonium Hydroxide ________________
Writing formulas for compounds with metals
with more than one oxidations state:
(Roman Numerals will be given)
1.
If there are Roman Numerals with the name, the Roman Numeral is
the charge of the metals.
For example: Iron (II) Oxide= Fe+2O-2=Fe2O2=FeO
2. If the compound ends in –ate or –ite, most likely you should look on
the polyatomic ion chart.
3. Write the metal, then look up the polyatomic ion and place it in
parenthesis. Put the charges on top and criss-cross. Reduce if
necessary.
For example:
Zinc Carbonate=
Gold (III) Thiocyanate =
For each compound listed below, write
the correct formula using the stock
system.
1.
Iron (II) Chloride _____________________
2.
Lead (IV) Phosphide __________________
3.
Tin (II) Oxide ________________________
4.
Copper (I) Iodide _____________________
5.
Nickel (III) Sulfide ____________________
6. Cobalt (II) Thiocyanate ______________
7. Manganese (IV) Oxide _________________
8. Titanium (IV) Chromate _________________
9. Iron (III) Sulfate _______________________
10. Lead (II) Nitrate ______________________
11. Tin (IV) Carbonate ______________
12. Copper (II) Acetate _____________
Writing Formulas for Covalent
Compounds
• Sometimes the name will have prefixes
(only when the two elements in the
compound are nonmetals). Simply use the
prefixes to figure out the formula!
For each compound listed below, write
the correct formula using the prefix
system.
1.
Diphosphorous pentoxide _____________________
2.
Silicon tetrafluoride ________________________
3.
Dihydrogen monoxide ________________________
4.
Tetraphosphorous trisulfide __________________
Naming Ionic Compounds using the Formula
Determining subscripts of
elements or polyatomic in
ionic compounds
Use Criss-cross method
Example:
Naming ionic compounds with
metals with only 1 oxidation
state
Binary Ionic Compounds
Ionic compounds with only 2 different elements.
For example: CaBr2 = Calcium
Bromide
Name the metal and end the nonmetal in –ide.
Ternary Ionic Compounds
Ionic Compounds with 3 different elements.
•Name the metal and then name the polyatomic ion.
Naming ionic compounds with
metals with multiple
oxidation state
1. Determine the oxidation state of the metal in the
compound.
2. Name the metal, put the oxidation state in ROMAN
NUMERALS in parenthesis and end the nonmetal in
–ide.
For example: KNO3 = Potassium
Nitrate
Example: FeO= Iron
(II) Oxide
Writing the Formula of Ionic Formulas using the name of the
Ionic Compound
Formula Writing for Ionic
Compounds
Writing formulas for compounds
with metals with more than one
oxidations state:
(Roman Numerals will be given)
All ionic compounds must be in empirical
form. (Reduced Form)
If the compound ends in –ide, most likely it
is a binary compound. Except for cyanide
and hydroxide.
If the substance ends in –ate or –ite it
contains a polyatomic ion. Put the polyatomic
ion in parenthesis before you criss-cross.
Put the oxidation states on top and crisscross the numbers.
If there are Roman Numerals with the
name, the Roman Numeral is the charge of
the metals.
2. If the compound ends in –ate or –ite,
most likely you should look on the
polyatomic ion chart.
3. Write the metal, then look up the
polyatomic ion and place it in parenthesis.
Put the charges on top and criss cross.
Reduce if necessary.
For example:
Zinc Carbonate= Zn+2(CO3) -2 =
Zn2(CO3)2 =Zn(CO3)
Gold (III) Thiocyanate = Au+3 (SCN)
Au(SCN)3
-1
=
For example:
Iron (II) Oxide= Fe+2O-2=Fe2O2=FeO
THE END
Naming Compounds
• The POLYATOMIC ION CHART (TABLE
E)
• These are ions consisting of more than
one atom. These names of polyatomic
ions ens in –ate or –ite, except for two (
ammonium and cyanide).
Naming Ionic Compounds
• 1. Cation first, then anion
• 2. Monatomic cation = name of
the element
• Ca2+ = calcium ion
• 3. Monatomic anion = root +
-ide
• Cl = chloride
Naming Ionic Compounds
(continued)
Metals with multiple oxidation states
• - some metal forms more than
one cation
• - use Roman numeral in name
• PbCl2
• Pb2+ is cation
How to name compounds:
• 1. Determine what the elements are in the
compound.
– A. If the compound is made of a metal and a
nonmetal, name as follows:
• 1. Name the metal first.
– IF THE METAL IS S TRANSITION METAL WITH MORE THAN
ONE OXIDATION STATE, TEHN YOU MUST ADD A ROMAN
NUMBERAL INDICATING THE OXIDATION STATE AFTER
THE METAL NAME.
– 2. Name the NONMETAL second but change the ending to –
ide.
– 3. Do not worry about how many of each element there are.
This does not matter in the name, just give the names of the
elements.
• 1. NaCl ____________________
• 2. K2O _____________________
• 3. MgBr2 ____________________
•
• 4. Al2S3 _____________________
• 5. Fe2O3 ____________________
• 6. FeO ______________________
7. AuCl3 _____________________
8. LiI ________________________
b. If the compound is made up of more than 2
different elements, check to see if there is a
POLYATOMIC ION in the compound.
Refer to the polyatomic chart. You should
familiarize yourself with these ions so that
you can easily recognize them.
1. If the compound is composed of a metal
and a negative polyatomic ion:
Metal + negative polyatomic ion
- name the metal
- then name the polyatomic ion.
MgCO3 _____________________________
If the metal is a transition metal with more
than one possible oxidation state, you
must indicate the oxidation state using a
Roman Numeral in parenthesis after the
name of the metal.
Ni(NO3)3 _______________________________________
2. If the compound is between a positive
polyatomic ion and a nonmetal:
polyatomic ion + nonmetal
- name the polyatomic ion first
- then name the nonmetal and change
the ending to –ide.
NH4Cl _____________________________
3. If the compound is between two
polyatomic ions: polyatomic ion +
polyatomic ion
- name the positive ion first
- name the negative ion second.
• NEVER CHANGE THE ENDINGS OF
POLYATOMIC IONS!!!!!!!
Practice:
1. NaNO3 ________________________
________________________
• 2. Al2(SO4)3 _______________________
_____________________________________
• 3. Ca(NO3)2 _______________________
•
_______________________
4. NH4Br ___________________________
___________________________
5. CuSO4 ___________________________
___________________________
c. If the compound is made of 2 nonmetals,
there are two ways to name it.
You can name the compounds the same way
you name the metals and nonmetals but
you must use an oxidation number roman
numeral in parenthesis to indicate the
oxidation number of the first element.
Naming Binary Compounds
Compounds between two nonmetals
- First element in the formula is
named first.
- Second element is named as if it
were an anion.
- Use prefixes
- Only use mono on second element -
• -
•
•
•
•
P2O5 = diphosphorus pentoxide
CO2 = carbon dioxide
• 1. CO2 _____________________
• 2. CO ______________________
• 3. P2O5 ____________________
• 4. SiO2 _____________________
• 5. N2O4 ____________________
Naming Acids
1. Elements that begin with hydrogen are
normally acids.
2. Naming Binary Acids
a. Begin with the word HYDRO- and end with
the non metal in –IC
b. Then add the word ACID
For example: HCl= Hydrochloric acid
3. Naming Ternary Acids
a. Cover the hydrogen and look up the
polyatomic ion being used.
b. If the polyatomic ion ends in –ITE change
the ending to –OUS and add the word ACID.
For example: HClO2 = chlorite = Chlorous
Acid
c. If the polyatomic ion ends in –ATE change
the ending to –IC and add the word acid.
For example: H SO = sulfate = Sulfuric Acid
Name the following acids:
1. HCl _______________________________
2. HBr _______________________________
3. HF _______________________________
4. H2S ______________________________
5. H3P ______________________________
11. H3PO4
_________________________
12. H2SO3
_________________________
13. H2S2O3
________________________
6. H2Se ___________________________
7. HI ______________________________
8. H2CO3 __________________________
9. H2C2O7 _________________________
10. HNO2 __________________________
For each of the following acids
below, write the correct formula.
1. Sulfuric Acid ________________________
2. Hydrochloric Acid ____________________
3. Clorous Acid ________________________
4. Cloric Acid __________________________
5. Thiocyanic Acid
_______________________
6. Acetic Acid _______________________
7. Nitric Acid
___________________________
8. Dichromic Acid _______________________
9. HydroIodic Acid ______________________
10. Nitrous Acid ________________________
How can you tell the difference between a binary
acid and a ternary acid from its name?