1. SYMBOLS
2. ATOMIC NUMBER AND MASS
3. SUBATOMIC PARTICLES
4. NOMENCLATURE AND FORMULA WRITING
5. MOLAR MASS
6. MOLE
THE PERIODIC TABLE
PERIODS
GROUPS
Atomic Numbers and Masses
• The atomic number of an
element is the number of
protons in the nucleus of an
atom of that element.
• It is the number of protons
that determines the identity
of an element
• Because atoms have no
overall electrical charge, an
atom must have as many
electrons as there are protons
in its nucleus.
• The sum of the protons and
neutrons in the nucleus is the mass
number of that particular atom.
• Isotopes of an element have different
mass numbers because they have
different numbers of neutrons, but
they all have the same atomic number.
37
• 35
𝐶𝑙
and
17
17𝐶𝑙
• In order to have a simpler way of
comparing the masses of individual
atoms, chemists have devised a
different unit of mass called an atomic
mass unit, which is given the
symbol u.
BRIEFLY,
Average atomic mass is
the average mass of all
the isotopes of an
element weighted by
their abundance.
1 THE SYMBOL (the name of the atom may
be written in the exam paper)
2 THE ATOMIC NUMBER (It is the number
of protons of the nucleus of the element
and also the number of electrons if the
atom is neutrally charged)
3 AVERAGE ATOMIC MASS (measured in
atomic mass unit (amu). It also
represents the average mass for a mole
of an element measured in grams. 1
mole of carbon atom has a mass of12.01
g. This is called molar mass of the
element.
=(36.966 x 0.242) + (34.969x 0.758)
= 8.946 + 26.507
= 35. 453
• Groups are numbered in two ways : Roman numerals (old type) and just
simple numbers. No need to memorise the the group numbers but some
groups are important to know. A groups are represantative , B groups are
transition elements.
Group 1 (IA) Alkaline metals
Group 2 (II A) Alkaline earth metals
Group 3-12 (Group B) Transition metals
Group 17 (VII A) Halogens
Group 18 (VIII A) Noble gases
The bottom rows are lantanides and actinites. Names are given according to
the first element. The other name of these groups are inner transition metals.
SUBATOMIC PARTICLES
• Electrons are a type of subatomic particle with a negative charge.
• Protons are a type of subatomic particle with a positive charge. Protons
are bound together in an atom's nucleus as a result of the strong nuclear
force.
• Neutrons are a type of subatomic particle with no charge (they are
neutral). Like protons, neutrons are bound into the atom's nucleus as a
result of the strong nuclear force.
• Protons and neutrons have approximately the same mass, but they are
both much more massive than electrons (1836 times as massive as an
electron).
• The positive charge on a proton is equal in magnitude to the negative
charge on an electron. As a result, a neutral atom must have an equal
number of protons and electrons.
• The atomic mass unit (amu) is a unit of mass equal to one-twelfth the
mass of a carbon-12 atom
Elemental Names and Symbols
• There are about 92 naturally occurring elements known on
Earth. Using technology, scientists have been able to create
nearly 30 additional elements that do not occur in nature. Today,
chemistry recognizes 118 elements.
• Today, new elements are usually named after famous scientists.
• The first letter of the symbol is usually the first letter of the
element’s name, while the second letter is some other letter
from the name. Some elements have symbols that derive from
earlier, mostly Latin names, so the symbols may not contain any
letters from the English name.
• In the universe as a whole, the most common element is
hydrogen (about 90% of atoms), followed by helium (most of the
remaining 10%). All other elements are present in relatively
minuscule amounts, as far as we can detect.
• On the planet Earth, however, the situation is rather different.
• Oxygen makes up 46.1% of the mass of Earth’s crust (the
relatively thin layer of rock forming Earth’s surface), mostly in
combination with other elements, while silicon makes up 28.2%.
Hydrogen, the most abundant element in the universe, makes
up only 0.14% of Earth’s crust.
Chemical Formulas
• A chemical formula is an expression that shows each of the
elements in a compound and the relative proportions of
those elements. Water is composed of hydrogen and oxygen
in a 2:1 ratio and its chemical formula is. Sucrose (table
sugar) consists of carbon, hydrogen, and oxygen in a
12:22:11 ratio.
• Oxygen and sulfur in water and sulfuric acid, respectively, do not
have a "1" subscripts
• Sometimes certain groups of atoms are bonded together within the
chemical and act as a single unit. Polyatomic ions are enclosed in
parenthesis followed by a subscript if more than one of the same ion
exist in a chemical formula. For example, the formula represents a
compound with:
Ca3(PO4)2
• Ca 3 atoms and
2 polyatomic ions (PO4)
• To count the total number of atoms for formulas with
polyatomic ions enclosed in parenthesis, use the subscript as
a multiplier for each atom or number of atoms.
NOMENCLATURE OVERVIEW
Binary Compounds
• Binary compounds are compounds that consist of only two elements.
Some binary compounds have special names, and these special
names supersede any of the rules given below.
• H2O is water, NH3 is ammonia, and CH4 is methane.
-
• All other binary compounds have a name with a suffix ide. Binary
compounds may be subdivided into metal type, nonmetal type, and
acid type.
METAL TYPE
• These binary compounds begin with metals. The metal is
given first in the formula. First name the metal, then name
the nonmetal with the suffix -ide.
• The ammonium ion (NH4+) is often treated as a metal, and
its compounds are named under this rule.
NH4Cl
NH4+ , Cl-
ammonium chloride
NONMETAL TYPE
• These binary compounds have formulas that begin
with a nonmetal. PREFIXES are used to indicate the
number of each atom present.
• No prefixes are used for hydrogen.
 In normal nomenclature, the nonmetal
prefixes are not used if a metal is present.
 One of the few exceptions to this is
MnO2, sometimes called manganese
dioxide.
ACID TYPE
• These binary compounds have formulas that begin
with hydrogen.
• If the compound is not in solution, the naming is
similar to that of the metal type.
• If the compound is dissolved in H2O, indicated by
(aq), the compound takes on the prefix hydro and
the suffix -ic.
• If the compound is not in solution, the state of
matter should be shown
HCl (g),HF (l)
• If the formula has no designation of phase or water,
either name may be used.
Ternary compounds
• Ternary compounds are those containing three or more
elements.
• If the first element in the formula is hydrogen, it is
usually classified as an acid.
• If the formula contains oxygen in addition to the
hydrogen, the compound is usually classified as an
oxyacid.
• If the first element in the formula is not hydrogen, the
compound is classified as a salt.
• Ternary acids are usually named with the suffixes -ic or ous. These acids undergo many reactions to form salts,
compounds of a metal, and the ion of an acid.
• If an acid name has the suffix -ic , the ion of this acid has a
name with the suffix -ate.
• If an acid name has the suffix -ous, the ion has a name
with the suffix -ite. Salts have the same suffixes as the
suffixes of the ions.
• When the ternary compound is not an acid, the first
element is usually a metal. In these cases, the name of
the compound is simply the name of the metal
followed by the name of the ion.
• The ammonium ion is treated as a metal in these
cases.
Writing Formulas
• To write the formula from the name of a binary compound containing
ONLY NONMETALS, write the symbols for the separate atoms with
the prefixes converted to subscripts.
• In all compounds, the total charge must be zero. There are NO
exceptions.
• It is necessary to have some idea what the individual
charges are. The species with the positive charge is listed
and named first; this is followed by the species with the
negative charge.
• Subscripts may be needed to make sure the sum of the
charges (valances) will equal zero.
• If a polyatomic ion must be increased to achieve zero
charge, parentheses should be used.
• For practice purposes it will be referred to as THE
CRISSCROSS RULE. It works most of the time and
therefore is worth considering.
• If the crisscross rule is applied, you should reduce the
formula if possible.
CHARGES ACROSS A PERIOD
 Metalloids may be treated as metals or nonmetals
Transition Metals
• Many transition metals and the group of six elements centered around
lead on the periodic table commonly have more than one valence.
• The valence of these metals in a compound must be known before the
compound can be named. Modern nomenclature rules indicate the
valence of one of these metals with a Roman numeral suffix (Stock
notation).
• Older nomenclature rules used different suffixes to indicate the charge.
• Stock notation is often omitted for Zn, Cd, and Ag, as they
do not have variable valences.
MOLAR MASS (M)
The mass (in grams) of one mole of an element or
compound
• just add up the masses of all the elements In the
compound (number of atoms x molar mass of the
element).
• Molar mass of H2O = 18.0 g (18.0 g mol-1)
Molar mass of C6H12O6 = 180 g (180 g mol-1)
THE MOLE AND EMPIRICAL FORMULA
MOLE
The amount of a substance that contains the same
number of particles as the number of atoms in 12 g
of carbon-12
• SI unit for amount of a substance
AVOGADRO’S NUMBER
The number of particles in one mole of a pure substance = 6.022 x
1023
• This can mean 6.022 x 1023 Atoms (elements), molecules (covalent
compounds), formula units (ionic compounds)
• 1 mole of C-12 = 6.022 x 1023 atoms of C-12 = 12.0 g of C-12
• 1 mole of glucose = 6.022 x 1023 molecules of C6H12O6 = 180.0 g of
C6H12O6
• 1 mole of NaCl = 6.022 x 1023 formula units of NaCl = 58.5 g of NaCl
MOLECULAR MASS
The mass of one mole of a molecule
• Molecular mass of H2O = 18.0 g
Molecular mass of C6H12O6 = 180 g
RELATIVE MOLECULAR MASS / RELATIVE FORMULA
MASS (Mr)
The mass of an average molecule of a compound relative
to 1/12 of the mass of an atom of carbon-12.
• The term relative molecular mass is used for covalent
compounds; for ionic compounds, the term relative
formula mass is used.
• The units of molar mass are g mol-1, but relative molar
mass has no units: the Mr of glucose (C6H12O6) is
180.18.
water
ofofwater
THE
NUMBER OF
PARTICLES
BALANCING EQUATIONS
204 g
204 g
EXAMPLE
a- What is the ratio of aluminium to oxygen atoms in aluminium nitrate
(Al(NO₃)₃)?
1mol Al(NO₃)₃ molecule consists of 1 mol Al atom, 9 mols of O atom. 1:9
b- Given the following information, a 1.00 g sample of which of the
following substances contains the greatest amount of chlorine?
LiCl, NaCl, KCl, RbCl 7 g of Li and 35.5 g Cl
23 g Na and 35.5 g Cl
39 g K and 35.5 g Cl
85.5 k Rb and 35.5 g Cl
FORMULAS
1. PERCENT COMPOSITION
% By mass of each element in a compound
EXAMPLE
Determine the percentage of magnesium present in magnesium
phosphate.
Molar mass of Mg3(PO4)2 is
M = (3 × 24.31) + (2 × 30.97) + (8 × 16.00) = 262.87 g mol–1
% Mg = (3 × 24.31)/ 262.87× 100 ≈ 27.7 %
EXAMPLE
Calculate the mass of oxygen present in 2.20 g of CO2. (1.60 g)
EXAMPLE: What mass of HNO3 contains 2.00 g of oxygen? (2.63 g)
2. SIMPLEST (EMPIRICAL) FORMULA
Formula showing the smallest whole number ratio of atoms in a
compound.
3. MOLECULAR FORMULA
The simplest (empirical) formula may not be the correct
formula!
For example, most sugars have the simplest formula CH2O;
but glucose is C6H12O6 (same ratio, but not exact same
formula:
CH2O Is the sImplest formula; C6H12O6 Is the molecular
formula
DETERMINE THE EMPIRICAL FORMULA FROM THE
PERCENTAGE COMPOSITION
EXAMPLE
a- You have a compound composed of boron & hydrogen.
The compound is 78.0% boron & 22.0 % hydrogen find the
empirical formula. (B: 10.81g/mol, H: 1.01 g/mol)
• Change % to grams
• Convert grams to moles
• Divide both numbers by the smallest number
• 𝐵 𝑚𝑜𝑙
• H 𝑚𝑜𝑙
78.0
=
= 7.22
10.81
22.0
=
= 21.8
1.01
7.22
•𝐵 =
= 1.00
7.22
21.8
• H=
= 3.02 round to 3
7.22
b) After an experiment you discover
that the formula mass of the true
compound is 27.67 u. What is its
molecular formula?
• Empirical formula BH3
X2
EXAMPLE
A carbohydrate, which contains C, H, and O, has a % composition of
33.3% C and 7.4% H. Find the empirical formula of this
carbohydrate. (O is 59.3 %)
C: 33.3 g / 12.01 g = 2.773 moles C
H: 7.4 g / 1.008 g = 7.34 moles H
O: 59.3 g / 16.00 g = 3.70625 moles
C: 2.773 / 2.773 = 1
H: 7.4 / 2.773 = 2.66
O: 3.70625 / 2.773 = 1.33
Multiply all of them by 3 to get whole
numbers.
C: 1 x 3 = 3 atoms C
H: 2.66 x 3 = 8 atoms H
O: 1.33 x 3 = 4 atoms O
Empirical formula C3H8O4
COMPOSITION BY MASS FROM COMBUSTION DATA
EXAMPLE :
An organic compound, A, contains only carbon and hydrogen.When
2.50 g of A burns in excess oxygen, 8.08 g of carbon dioxide and
2.64 g of water are formed. Calculate the empirical formula.
𝑦
𝑦
CxHy + (x + ) O2 → xCO2 + H2O
4
2
number of moles of CO2 = 8.08/ 44.01 = 0.184 mol
number of moles of H2O= 2.64/ 18.02 =0.147mol
Multiply
by
the
smallest number to
get a whole number
for both atoms
X5
C5 H8
REACTING MASSES AND VOLUMES
ConservatIon of mass (lavoıser)
• If 55.85 g of iron reacts exactly and completely with 32.06
g of sulfur, 87.91 g of iron sulfide is formed:
Fe(s) + S(s) → FeS(s)
UsIng moles
• Consider the reaction of sodium with oxygen:
4Na(s) + O2(g) → 2Na2O(s)
A how much sodium reacts exactly with 3.20 g of oxygen?
B what mass of Na2O is produced?
a Number of moles of oxygen= 3.20/32.00
=0.100 mol
1mol O2 reacts with 4mol sodium. 0.100 mol O2
reacts with
4 × 0.100 mol = 0.400 mol sodium.
0.400×22.99=9.20g Na
B 1mol O2 yields 2 mol Na2O, 0.100 mol yields
-1
0.200 mol Na2O 0.200 mol x 61.98 gmol = 12.4
g
If one of the reactants is consumed, the other is in excess quantity
EXAMPLE:
SiO2(s) + 4HF(g)
SiF4(g) + 2H2O(l)
If 2.0 mol of HF is combined with 4.5 mol SiO2
A)which is the limiting reactant? HF
B) what are the moles of the limiting reagent and the
excess reagent during the reaction?
0.50 mol SiO2 and 2.0 mol HF
C) what is the mol of the excess reagent after the
experiment?
4.5 mol- 0.50 mol = 4.0 mol SiO2
EXAMPLE:
C6H6(l) + Cl2(g)
C6H5Cl(s) + HCl(g)
• When 36.8 g of C6H6 react with excess Cl2, 38.8 g of C6H5Cl(s) are
produced (the actual yield). What is the percent yield of C6H5Cl(s)?
36.80
= 0.4714 mol C6H6
78.06
38.80
= 0.3449 mol C6H5Cl
112.50
1:1 mol in reaction
0.3449
x 100 = 73.17 %
0.4714
EXAMPLE
2 HBr(aq) + Zn(s)
ZnBr2 (aq) + H2(g)
A piece of solid zinc weighing 98 grams was added, to a solution
containing 324 grams of HBr. What is the volume of H2 pressure at STP if
the reaction above runs to completion?
98 𝑔
nZn =
= 1.5 mol
65.38 𝑔/𝑚𝑜𝑙
324 𝑔
nHBr =
= 4 mol
80.98 𝑔/𝑚𝑜𝑙
According to the reaction Zn is the limiting reactant.
HBr: Zn will be 3: 1.5.
1.5 mols of H2 gas must be produced.
At STP the volume of H2 will be 33.6 L.