Lecture Notes for Inorganic Chemistry

CHEM 1 [GENERAL & INORGANIC CHEMISTRY]
Aklan State University-Ibajay Campus |1st Semester, 2015-2016
© Kenneth D. Barrientos
1
BASIC ATOMIC STRUCTURE
(Atoms, Molecules, & Ions)
THREE FUNDAMENTAL CHEMICAL LAWS
A. LAW OF CONSERVATION OF MASS (Antoine Lavoisier, 1743 – 1794)
 Matter is neither created nor destroyed.
Example:
2H2 + O2 → 2H2O
4H + 2O 4H + 2O =
B. LAW OF DEFINITE PROPORTIONS
 A given compound always contains exactly the same proportion (ratio) of elements by mass.
Example: Water, H2O
2 H = 2 (1.01) = 2.02
1O = 1 (16.00)= 16.00
----------------
18.02
 In other words, the ratio of H:O atoms in water is always 2H:1O within the one, definite
compound H2O.
When John Dalton was on his way to discovering the “atom”, he stumbled into, and thus discovered
the next law…
C. LAW OF MULTIPLE PROPORTIONS
 When two elements form a series of compounds, the ratios of the 2nd elements that combine
with the first element can be reduced to small whole numbers.
Example: CO vs. CO2
 In the multiple compounds, the ratio of “O” atoms per the one “C” atom is 2:1.
 Unlike the Law of Definite Proportions, here we are considering multiple compounds and the
ratio is between two of the same type of element.
Dalton began to think that because there was always whole-number multiples, there must be some
simple “unit”. This he called an atom.
Same number and type of
elements (matter) on either side
H2O is always 11.2% H
and 88.8% O by mass
Atomic Weight
For example, one cannot have CO1.5 because you cannot have 0.5 of a unit (an oxygen atom)

2
Dalton’s Atomic Theory (1805)
1. All matter consists of atoms.
2. All atoms of a given element are identical, having the same size, mass and chemical
properties. The atoms of one element are different from the atoms of all other elements.
3. Compounds are formed when atoms of different elements are combined. A given compound
always has the same relative numbers and types of atoms.
4. Chemical reactions only involve the rearrangement of atoms. Atoms are not created or
destroyed in chemical reactions.
5. Atoms cannot be divided any further
200 years later, Dalton’s Atomic Theory (above) only has two modifications that must be made:
 In theory 2 above, atoms of a given element are not always identical (isotopes)
 In theory 5, atoms can be divided further into the 3 subatomic particles: proton, electron and
neutron
EARLY EXPERIMENTS TOO CHARACTERIZE ATOMS
ELECTRONS (e-)
 Discovered in 1903 by J.J. Thomson as he conducted his experiment with the cathode ray
tubes. Cathode rays are streams of electrons running through the tube from cathode to
anode.
 Because the “rays” were deflected away from the negative end of an applied electric field,
Thomson postulated that the cathode rays were negatively-charged particles called electrons
(e-).
 Because atoms were known to be neutral, Thomson reasoned that there must be a positive
charge somewhere in the atom as well. Hence, he came up with the “Plum Pudding Model” of
the atom.

3
Robert Millikan (1909) – performed experiments that determined the charge of an electron
and its mass.
Mass of e- = 9.10 x 10-28 g or 9.10 x 10-31 kg
Charge of e- = -1.60 x 10-19 C
Ernest Rutherford (1911) – disapproved Thomson’s “Plum Pudding Model” of the atom and
discovered the nucleus. His experiment involved the bombarding of various atoms of metal
foil (gold foil) with alpha particles.
 Most alpha particles pass directly through because atom is mostly open space.
 Deflected particles are those that had some “close to” the positively-charged center.
 Reflected particles had a “direct hit” with the center.
 The densely-packed center of an atom is called nucleus and represents 99.99% of an
atom’s mass.

4
ATOMS
 Built from elementary particle.
 Atoms cannot be sub-divided by chemical reactions: available energies are not high enough
 Atoms can be subdivided, modified or transferred into new atoms by physical reactions
(nuclear reactions) Reason: energies are sufficiently high
 Chemical properties of atoms are significantly determined by the elemental building units of
atoms
 The electron’s mass is 10,000 times smaller than the mass of proton or neutron.
 a.m.u. = atomic mass unit = 1/12 the mass of a carbon-12 atom.
 Question: If all atoms are composed of electrons, protons and neutrons, why do different
atoms have different properties?

5
PROTONS and ELECTRONS
 The number of protons or electrons in an atom or ion determines what element it is.
 For example, if a particle has 6 protons in it, it must be CARBON.
ATOMIC NUMBER (Z)
 The atomic number of an atom or an ion is equal to its number of protons/electrons.
Atomic Number = Number of Protons or electrons
Example:
Element # of protons # of electrons Atomic # (Z)
Carbon 6 6 6
Phosphorus 15 15 15
Gold 79 79 79
ATOMIC MASS NUMBER (a.m.u)
 The atomic mass number of an atom or ion is equal to the sum of its number of protons plus its
number of neutrons.
A.M.U. = # protons + # neutrons
Example:
Nuclide p+ n0 e- Mass #
Oxygen 8 10 8 18
Arsenic 33 42 33 75
Phosphorus 15 16 15 31
IONS
 Are charged particles
 Ions are formed when atoms gain or lose electrons.
 Positive ions (cations) are formed when a neutral atom loses electrons.
 Negative ions (anions) are formed when a neutral atom gains electrons.
 Metallic atoms tend to lose electrons to form positive ions (also known as cations).
 Nonmetallic atoms tend to gain electrons to form negative ions (a.k.a. anions)
CHARGE (NEUTRON – n0
)
 The charge of an atom or ion is equal to its number of protons minus its number of electrons.
Charge = # protons - # electrons
Example: A particle with 34 protons and 36 electrons has a charge of -2.

6
ISOTOPES
 Isotopes are atoms of the same element that have a different number of neutrons. Therefore,
isotopes have the following characteristics:
1. Isotopes have the same atomic number (same number of protons), but a different atomic
mass number (a different number of neutrons).
2. Isotopes behave the same chemically, because they are the same element. The only
difference is that one is heavier than the other, because of the additional neutrons.
For example: Carbon-12 and Carbon-14 are both isotopes of carbon. Carbon-12 has 6
neutrons; carbon-14 has 8 neutrons.
 Isotopes can be represented in two notations:
1. Nuclear Symbol Notation 2. Hyphen Notation

7
ATOMIC WEIGHT
 The atomic weight of an element (as it appears in the periodic table) is the weighted average
of the atomic mass numbers of all of the isotopes for that element.
 For example, the atomic weight of carbon (as shown on the periodic table) is 12.011 a.m.u.
(atomic mass units). However, there is no one carbon atom that has a mass of 12.011 a.m.u.
Carbon exists as four different isotopes: carbon-11, carbon-12, carbon-13, and carbon-14,
which have approximate atomic mass numbers of 11, 12, 13, and 14 a.m.u., respectively. If
you know the percent abundance of each of those isotopes, you can calculate the atomic
weight of carbon by determining the weighted average of the atomic mass numbers of the
four isotopes. That weighted average comes to 12.011 a.m.u. The reason that the atomic
weight is closer to 12 than it is to the other atomic mass numbers is that carbon-12 is the most
common isotope of carbon.
 For most elements, the atomic mass number of the most common isotope for that element
can be determined by rounding the atomic weight for that element to the nearest whole
number. For example, the atomic weight of aluminum is 26.98154 a.m.u. Therefore, the most
common isotope for aluminum can be assumed to be aluminum-27.
ISOTOPE Symbol Composition of the nucleus % in nature
Carbon-11 11C
6 protons
5 neutrons
<0.01 %
Carbon-12 12C
6 protons
6 neutrons
98.89%
Carbon-13 13C
6 protons
7 neutrons
1.11%
Carbon-14 14C
6 protons
8 neutrons
<0.01%

8
MOLECULES AND IONS
 The forces that hold two or more atoms together are called chemical bonds.
1. COVALENT BONDS – when two atoms share electrons to form a molecule.
Example:
Methane, CH4 →
Chemical formula
2. IONIC BONDS – when two atoms attract each other electron to form ions.
 Cations – positively charged ions
 Anions – negatively charged ions
Example: Na+1 + Cl-1 → NaCl
Practice Yourself:
1. For each of the following atoms or ions, provide the number of protons, neutrons, and
electrons:
a) 40Ar
b) 40Ca2+
c) 39K+
d) 39K
2. Calcium chloride is an ionic substance. Determine the number of protons and electrons in a
calcium ion (as found in calcium chloride).
3. The chloride ion exists in two common isotopes: 35
Cl–
and 37
Cl–
. Determine the number of
protons, neutrons, and electrons for both isotopes.
Structural Formula

9
INTRODUCTION TO PERIODIC TABLE
The modern periodic table was developed by the Russian chemist Dmitri Mendeleev (1834-1907) and
the German chemist Lothar Meyer. Around 1869 Mendeleev recognized the periodicity of the
elements. Although scientists previously classified the elements, Mendeleev succeeded in interpreting
the principles in an adequate way.
Periodic Table
 Shows all known elements in the universe.
 Organizes the elements by chemical properties.
 Mendeleev realized that after a certain number of elements similar or related properties
appeared again.
 He ordered these elements vertically underneath each other, but still ordering the elements
according to their atomic mass.
 Criteria: (a) increasing atomic mass, (b) chemically similar elements below each other
What is an ELEMENT?
o A substance composed of a single kind of atom.
o Cannot be broken down into another substance by chemical or physical means.
Structure of Elements
o Monatomic elements composed of individual discrete (atoms) He-1
o Diatomic elements composed of molecules made up of two atoms that are bonded together
Br2, I2, N2, Cl2, H2, O2, F2 (Brinclhof).
What is a COMPOUND?
 A substance in which two or more different elements are CHEMICALLY bonded together

10
What is a MIXTURE?
o Two or more substances that are mixed together but are NOT chemically bonded.
Types of Elements
a. METALS
 Can be found in the middle and on the left off the Periodic Table.
 Most elements are metals
 Metals are good conductors of heat and electricity.
 Metals are shiny.
 Metals are ductile (can be stretched into thin wires).
 Metals are malleable (can be pounded into thin sheets).
 A chemical property of metal is its reaction with water which results in corrosion (lustrous).
b. NON-METALS
 With the exception f Hydrogen (a non-metal) these are found on the upper-right and right
side of the Periodic Table
 Chemically, no-metals tend t gain e-s to form anions.
 Non-metals are poor conductors of heat and electricity.
 Non-metals are not ductile or malleable.
 Solid non-metals are brittle and break easily.
 They are dull.
 Many non-metals are gases.
c. METALLOIDS (Semi-Metals)
 Share properties of both metals and non-metals.
 They are solids that can be shiny or dull.
 They conduct heat and electricity better than non-metals but not as well as metals (semi-
conductors).
 They are ductile and malleable.
 There are eight (8) of them, located in a staircase pattern between the metals ad
nonmetals:
Boron (B) Antimony (Sb)
Silicon (Si) Tellurium (Te)
Germanium (Ge) Polonium (Po)
Arsenic (As) Astatine (At)

11
STRUCTURE OF PERIODIC TABLE
Main Group Elements (Vertical Groups)
o Group 1(IA) - Alkali Metals
o Group 2(IIA) - Alkaline Earth Metals
o Group 13(IIIA) - Boron Family
o Group 14(IVA) - Carbon Family
o Group 15(VA) - Nitrogen Family
o Group 16(VIA) - Oxygen Family (Chalcogens)
o Group 17(VIIA) - Halogens
o Group 18(VIIIA) - Noble Gases
Other Groups (Vertical and Horizontal Groups)
o Group 3-12(IB - 8B) - Transition Metals
o Period 6 Group - Lanthanides (Rare Earth Elements)
o Period 7 Group – Actinides

12
NOTES
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Lecture Notes for Inorganic Chemistry

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