General Chemistry 1
Chemistry 16 Lecture
GENERAL CHEMISTRY 1
Chemistry 16
Science: dubbed as “body of knowledge”
✔ Computational Chemistry – facilitating with drugs
✔ Biochemistry
✔ Biotechnological Research
Scientific Method: cut and dried approach
1. Observation
Types of Data:
a. Qualitative: numerical facts
b. Quantitative: non numerical facts
Law: of various statements of observation; convenient way of storage of very large
amount of data. It allows us to predict yet an untried experiment.
It is a conclusive or explicit statement of fact that is evident to anyone
making the same observation.
1. Propose a tentative explanation.
Hypothesis: A tentative model.
Theory: describe a model that has been tested many times.
*If the hypothesis is incorrect, it is discarded or modified.
Definition of Problem
Collection of Data
LAW
Formulation of
Hypothesis
Testing Of Hypothesis
By Predicting Results
Of New Experiments
THEOR
Y
CHEMISTRY: is the study of matter, its composition, structure and properties, the
changes that it undergoes and the energy transformations accompanying these
changes.
Matter: anything that takes up space and has mass.
• Mass: a measure of its resistance to change in velocity
• Weight: measure to the force with which an object of a certain mass is
attracted by gravity to the earth or to some body near it.
Newton’s Equation:
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General Chemistry 1
Chemistry 16 Lecture
F = ma
W = mg
Force = Mass x Acceleration
Weight = Mass x Gravity
Composition: amount of components of matter; what is something is made of and
its quantities.
Structure: arrangement of particles that make up a substance.
Properties: Characteristics (Physical State, Mass, Color, Surface, Etc.)
A. Extensive: depend on size or amount of sample of matter
Ex.: mass, volume, length, heat
B. Intensive: independent of sample size
Ex.: physical state, color, hardness, melting point, density, specific gravity
Note: Density = Mass/Volume, is an intensive property (Ratio of two extensive
property)
A. Physical: observed without changing the chemical make-up of a substance.
a. Difficult to assign value – odor, taste, color
b. Can be expressed in definite numbers – melting, hardness
B. Chemical: Interaction between chemical substances.
Example: Fe + H20 = FeO2
Changes in Matter
A. Chemical Change: Results in disappearance of substances and formation of
new ones.
B. Physical Change: Does not result in formation of a new substance
Classification of Matter
A. Mixtures: consists of two or more substances and can be separated by
physical means.
Ex.: Blood, Air
Characteristics;
1. Consists of two or more substances
2. Has variable compositions
3. Can be separated by physical means
Types:
1. Homogeneous: also called solution. It has the same physical and
chemical properties through out; uniform
2. Heterogeneous: has physical and chemical properties that are not
uniform through out the sample; has two phases.
Ex.: Blood, Granite, Concrete
A. Pure Substances:
1. Have uniform properties
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General Chemistry 1
Chemistry 16 Lecture
2. Has definite composition
3. Cannot be further separated into other components by physical
means
1. Elements: Substances that cannot be made from or decomposed into
simpler substances; building blocks of more complex substances.
2. Compounds: Composed of two or more elements in fixed proportion by
mass; can be decomposed by chemical means.
Some Separation Techniques:
1. Thin Layer Chromatography (TLC)
FIGURE 1. Concept Map of Matter
MATTER
Mixture
PHYSICAL CHANGE
Variable Compositions
Pure Substances
Constant composition
Compounds
CHEMICAL CHANGE
Elements
MATTER
occurs as
Heterogeneous
Mixtures
SUBDIVISIONS
Homogeneous
Mixtures
Solutions
SEPARATION BY PHASE
Pure Substance
Chemical Elements
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General Chemistry 1
Chemistry 16 Lecture
CHEMICAL REACTIONS
Chemical Compounds
Energy: Capacity to do work.
An object can possess energy in just two ways
1. Kinetic Energy (KE)
2. Potential Energy (PE)
Total Energy of an object = KE + PE
A. Kinetic Energy: energy an object has when it is moving.
KE = 12mv2
Where: m = mass
v = velocity or speed
B. Potential Energy: Stored Energy
PE is not in used, but it is stored and has the capacity to do work when released
(Attractive or Repulsive)
An object attracted or repelled by some object has PE.
a)
Attracted
PE Increases
PE Decreases as the two balls come together
b)
Repelled
PE Increases
PE decreases
Some Forms of Energy
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General Chemistry 1
Chemistry 16 Lecture
Radiant Energy: Electromagnetic Radiations (Ex.: Infrared)
Atomic or Nuclear Energy: Manner atoms are built
Note: Amount of energy released or absorbed depends upon the amount of matter
allowed to react. (Intensive Property)
Energy is transformed. Final energy transformations may be light, sound, electricity
or heat.
SI (System International) unit for Energy:
Joules (1kg/sec2)
Heat
It is a form of energy that
flows by itself from high to
low.
Temperature
It is a measure of the
intensity of heat
Temperature Scales: 0K, 0C, 0F
SI Unit: Kelvin Scale (0K)
Calorie: Amount of heat needed to raise 10C of water.
Kilocalorie: Larger and more appropriate unit when dealing with chemical
reactions.
1 calorie
=
4.1840 J
1 kilocalorie =
4184.0 KJ
Law of Conservation of Energy: In an isolated system, like the universe, the total
energy is constant and energy is neither created nor destroyed but instead it can
only be transformed from one kind of energy to another.
Exothermic Reaction: Release of heat energy
Endothermic Reaction: Absorption of heat energy.
Law of Conservation of Mass: Mass is neither created nor destroyed in any
transformation of matter.
Law of Conservation of Mass – Energy:
E = mc2
where:
E = Energy
m = mass
c = speed of light
*The quantity of energy liberated or destroyed is exactly equal to the quantity of
matter destroyed or created.
Chemical Formula: shorthand way of writing names of compound. It has specified
composition of a complex chemical substance.
Atom: smallest particle of an element that can enter a chemical reaction.
Molecule: group of atoms that bond tightly together.
Chemical Equation: before and after picture of a chemical reaction.
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LAWS OF CHEMICAL COMBINATION
A. Law of Conservation of Mass: The total mass of the starting materials in a
reaction is the same as the mass of the product.
B. Law of Definite Proportion or Composition: by Joseph Proust.
In 1799, he shows that copper carbonate prepared in the laboratory had the
same component as naturally occurring copper carbonate.
Ratio of Cu:O:C = 5:4:1
Statement: For any sample of a pure chemical substance, its constitute element
exist in the same definite proportion.
Problem:
A sample of Silicon dioxide was found to contain 4.61 g silicon and 5.27 g oxygen.
How much oxygen would be combined with 12.0 g Si in another sample of Silicon
Dioxide?
Let x = amount of Oxygen combined with 12.0 g Si
x12.0 g Si= 5.27 g O4.61 g O
x=5.27 g O(12.0 g Si)(4.61 g O)
x=13.7 g O
ATOMIC THEORY
Idea of Atom:
Leucippus (Greek) - ~490 BC - ?
Democritus (Grek) ~470 – 380 BC
Atomos: meaning “indivisible”
Aristotle: Rejected the ides of atom. He believed that matter is continuous.
Epicurus: 341 – 270 BC
Lucretius: 95 – 55 BC, a Roman Poet. He wrote De Rerum Natura (On the Nature of
Things)
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Chemistry 16 Lecture
John Dalton’s Atomic Theory (English, 1803)
Postulates:
1. Matter is composed of indivisible particles called ATOMS.
2. All atoms of a given elements have identical properties.
3. Atoms of different elements have different properties.
4. A chemical reaction merely consists of reshuffling of atoms from one set of
combination to another. The individual atom themselves remains intact.
Atoms are not created, destroyed or change.
5. When atoms combine, they combine in fixed ratios of whole numbers forming
particles known as MOLECULES.
Billiard Ball Model: Dalton’s Atomic Model
– Tiny, hard, indestructible sphere.
A. Law of Multiple Proportions: When the mass of one element is the same in
two compounds, the masses of the second element are in a ratio of small whole
numbers.
To illustrate: NO, Nitric oxide and N2O, Nitrous oxide (laughing gas, for dental
anesthetic)
In NO, 0.875 g of N is present for every gram of O.
In N2O, 1.750 g of N is present for every gram of O.
0.875 g
1.750 g
=
1, a small whole number ratio of N to O.
2
Example 2: Sulfur forms 2 compounds with Oxygen.
Compound A
1.0 g S
1.0 g O
Compound B
1.0 g S
1.5 g O
Mass of Compound A
Mass of Compound B
1.0 g
O
1.5 g
O
2
x
2
2
=
3
ATOMIC STRUCTURE
Discovery that overturned Dalton’s concept of atom:
X-rays
1895
Radioactivity
1896
Electron
1897
Radium
1898
FIGURE 2. Parts of a discharge tube.
 JJ Thomson: devised discharge tube,
subdivisions of electrical charge
particle
 Michael Faraday
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General Chemistry 1
Chemistry 16 Lecture

Sir Humphry Davy
Properties of a Cathode Ray:
1. Consists of a stream of particles of definite mass.
2. Travel in a straight line away from the cathode.
3. Objects placed in their paths cast shadows on the end of discharge tube.
4. They are negatively charged, by the part that they are attracted to a
positively charged plate or attracted to a negative plate.
5. Nature of the cathode ray is the same irrespective of:
a. The material the cathode ray is made of.
b. Type of residual gas present in evacuated tube.
c. Kind of metal wire used to conduct electrical current.
d. The material used to produce current.
CHARGE TO MASS RATIO
JJ Thompson determine the charge to mass ratio of electron as
e
-1.76 x 108
=
coulumbs/gram
m
Where:
e = charge of electron in coulumbs
m = electron mass in grams
Robert Andrews Millikan: (1868 – 1953) devised oil – drop experiment.
In 1909, he determined charge of an electron
e = -1.60 x 10-19 c
m = 9.1 x 10-28 g
Charge of an ē (for convenience): -1
Proton: Positive Particle; was found to be massive than electron and dependent on
the kind of gas present in tube. It is the lightest, mass nearly equal to Hydrogen.
– 1.67 x 10-24, 1840 times of electron (Rest Mass)
– Convenient Charge of 1+
Henri Becquerel: Radioactivity
Components of Radiation
a. Beta (β) Particles: as electrons
b. Alpha (α) Particles: Charge of 2+ and mass of 7300x of electron
c. Gamma (γ) Rays: Magnetic Radiation like visible light but of considerably
high energy
TABLE 1. Components of Radiation
Name
Mass Relative to Hydrogen
Atom
Alpha (α)
~4
Beta (β)
1/1837
Gamma (γ)
0
Relative Charge
2+
10
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General Chemistry 1
Chemistry 16 Lecture
FIGURE 3. JJ Thomson’s Raisin Bread Model of
Atom
• Set of Positive Electrifications (+)
• Electrons (-)
Size of Nucleus: _______________
In 1908 to 1909, Ernest Rutherford, Hans Geiger and Ernest Marsden, used
gold in studying movement of atom and radioactivity.
Geiger counter: device used to count radioactivity
Nuclear Model of an Atom: formulated by Rutherford.
Mass Spectrograph: measure mass of atom.
In 1923, James Chadwick, discovered neutron (no charge). It has the mass same
as proton.
TABLE 2. Subatomic Particles
Particle
Charge
Coulomb
Charge Unit
+1.6022 x
+1
10-19
Location
Proton
Nucleus
Neutron
Electron
0
Around the
Nucleus
-1.6022 x 10
19
0
-
-1
Mass
Gram
1.6725 x 10
-24
1.6725 x 10
AMU
1.00728
-24
1.00867
9.109 x 10-28
0.000549
QUARKS AND LEPTONS
Charge
Up (u) Quark
Down (d) Quark
Protons
Neutrons
Electrons
1
- 3
/
2
+ 3
/
2 u quarks and 2 d quarks
1 u quarks and 2 d quarks
Structure less; belongs to a family of particle called LEPTONS
ATOMIC NUMBER: A concept of Henry Moseley.
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General Chemistry 1
Chemistry 16 Lecture
↓ Lower Wavelength (λ) Higher Atomic Mass (Inverse)
Shorter Wavelength (λ) Greater Energy of X-rays
✔ No. of positive charge – increases to another atom by one single electron unit.
(Atomic Number)
Atomic Number (Z): equal to the number of protons in the nucleus of an atom of a
particular element; also equal to the number of electrons


Element: a substance of all whose atoms contains the same number of protons.
Atom: extremely small, electrically neutral particle that has a tiny but massive
nucleus and one or more electrons relatively far from its nucleus.
Atomic Number = No. of Protons = No. of Electrons
Isotopes: Atoms of same element that has different atomic masses.
A nuclear symbol identifies an isotope of an element.
X = Atomic Symbol
Z = Atomic Number
A = Mass Number
Atomic Mass = Proton + Neutron
Atomic Mass = Atomic Number + Neutron
Atomic Mass = Electron + Neutron
816O - Oxygen Sixteen or O – 16
714N – Nitrogen – 14
715N – Nitrogen – 15
From name and symbol of each isotope, determine the number of nucleus.
Given: 612C
p=6
e=6
n=6
Consider 53131I, used in the treatment of cancer of thyroid gland and
hyperthyroidism.
A = 133
p = 53
n = 131 – 53 = 78
e = 53
e = 53
Nuclear Charge = 53 (Equal to
the No. of Protons)
Consider 3890Sr,
A = 90
p = 38
n = 90 – 38 = 52
e = 38
e = 38
the No. of Protons)
Nuclear Charge = 38 (Equal to
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ATOMIC MASS
Consider Carbon Monoxide (CO):
Mass O
Mass C
=
16 g O
12 g C
=
1.33
1
Consider Carbon Dioxide (CO2):
Mass O
Mass C
=
32 g O
12 g C
=
2.66
1
Relative Masses are called Atomic Masses
Reference: Carbon – 12, 12 amu
1 amu = 1/12 mass of C-12 atom.
Sample Problem: The average mass of Cu is 5.29 times greater than the mass of C12 atom. What is the atomic mass of Cu or Copper?
5.29
1
=
Cu Mass
Ave. mass of Cu atom
Ave. mass of C-12 atom
= 5.29 x 12 amu
= 63.5 amu
Sample Problem 2: An isotope of an element is ½ times as heavy as a C-12 atom.
Find the atomic mass of the element
At. Mass
= ½ x 12
= 6 amu
Given the isotopes of Mg and relative abundances, calculate the Average Atomic
Mass.
Isotop
e
Mg – 24
Mg – 25
Mg – 26
Isotopic
Mass
23.985
24.986
25.983
Isotopic Abundance
(%)
78.99%
10.00%
11.01%
Ave. Atomic Mass
Proportional Contribution to
Mass
18.95
2.499
2.860
24.309
Isotopic Mass x Relative Abundance = Proportional Contribution to Mass
Average Atomic Mass = ∑ Proportional Contribution to Mass
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NUCLEAR CHEMISTRY
Radioactivity: discovered in 1896 by Antoine Henri Becquerel (1852 – 1908), a
French Chemist. Ex.: Uranium
 Atoms of some elements are not stable
 They spontaneously disintegrate and emit radiation of various types –
RADIOACTIVITY.
 Radioactive isotopes or Radioisotopes: Change of Nuclei
 Transmutation: Change from one element to another (Rutherford and Sodi
Carnot)
Ex.: 88226Ra → 24He Alpha Particles+ 86222Rn
 Nuclear Reaction: involves a change in Atomic Number and/or the Mass
numbers
 In a typical radioactive decay reaction, mass numbers are conserved.
○ Charge is conserved. Thus, the sum of Z of the reacting nuclei and
the particles must be equal to the sum of the atomic numbers of the
product.
1327Al+24He → 01n+ 1530P
92240U→-10β+ 93243Np* (High Energy State)
93243Np*→γ+ 93243Np
Types of Radiation:
1. Alpha (α) radiation: composed of He2+ ions called alpha particles (α)
2. Beta (β) radiations: consists of electrons (β)
3. Gamma (γ) radiations: highly energetic, very penetrating light waves.
Small Particles involved in Nuclear Particles and Symbols
Alpha
24He or α
Electron or Beta
-10e or -10β
Positron (+e)
+10e or +10β
Proton (Hydrogen nuclei)
11H or 11p
Neutron
01n
Nuclear Stability: Zone or Belt of Stability
✔ Above: Unstable; too many neutron, spontaneous β- production
✔ Below: Unstable; too many protons, spontaneous positron production
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RADIOACTIVE DECAY
A. Decay of neutron rich nuclei
– Above the belt; excess of neutron and few protons
1. Β or Beta Decay: stability is achieved by emitting β particles.
01n → 11p+-10β
Ex.: 614C → 114N+-10β
n
C
=
8 =1.3
3
6
=
p
N
=
n
p
=
7
7
=1.0
0↓
3687Kr → 3787Rb+-10β
Kr
=
n
=
p
5
1
3
6
=1.4
2
Rb
=
n
=
p
5
0
3
7
=1.33
↓
2. Neutron Emission
3687Kr → 3686Kr+01n
Kr 87=
n
=
p
5
1
3
6
=1.4
2
Kr - 36
=
n
=
p
5
0
3
6
=1.39
↓
A. Decay Processes for Neutron poor Nuclei
1. Positron Emission: One less Proton and one more neutron
3578Br → 3478Se++10β
Br
=
n
=
p
4
3
3
5
=1.2
2
Se
=
n
=
p
4
4
3
4
=1.29
↑
2. Proton Emission: results in one less proton.
2143Sc → 2042Ca+11H
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3. Alpha Emission: Results in two less protons and two less neutrons
84208Po → 82204Pb+24He
4. Electron Capture or K – Capture: one less proton, one more neutron.
79195Au+e- → 78195Pb+γ
82205Pb+e- → 81205Pb+γ
Radiation Protection:
To prevent skin damage due to Beta particles, heavy clothing and wearing gloves is
encouraged.
For Gamma rays, lead or concrete can only prevent from emission. It is the most
hazardous among rays.
Radioactivity and Living Organisms
Cellular Damage by Radiation
1. Formation of ION PAIRS
H2O  H2O+ + e2. Formation of FREE RADICALS (unpaired electrons and quite unstable)
H2O  HO* + O*
Free radicals can:
1. Recombine to form H2O, which is harmless
2. Combine to form H2, tolerated in small amounts
3. Combine to form H2O2, which is highly toxic
4. React with oxygen in cell to produce a free radical that is even more
undesirable than H2O2
Half Life or Radioisotopes
Half Life or T½: time it takes ½ of a radioactive sample to decay
Sample: Consider 131I, with a half life of 8 days. If 200 mg of the radioisotope is left
to decomposed, how many mg will be left after 4 half life?
200 mg
100 mg
50 mg
25 mg
12.5 mg
0 day
8 days
16 days
24 days
32 days
1
2
3
4
Naturally Occurring Isotopes with long T ½
C – 14
K – 40
Half - life
5760 yrs.
1.3 x 109 yrs
Emitting Particles
β
β, γ
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Ra – 226
U – 238
1600 yrs
4.5 x 109 yrs
Medical Radioisotopes
C -11
20 mins
Cr -51
28 days
I – 131
8 days
K – 42
12 hrs
Sr – 85
64 days
Tc -99
8 hrs
Formula for Half – Life
lo N
0.30 T
g
1
0
=
N
T
½
α, γ
α, γ
β+
γ
β, γ
β, γ
γ
where: N0 = Initial Mass
N = Mass of Sample at Time t
T ½ = Half life
Sample: Se -75 has a T ½ of 120 days. If we begin at 8.00 g of Se-75, how may
grams would remain after 240 days?
Given: No = 8.00 g Se-75
T ½ = 120 days
t = 240 days
N=?
lo
g
N
0
N
log 8.0
0
=
0.30
1
T
T
½
0.30
= 1
N
log 8.0
0
N
8.0
(240
)
120
=
0.60
2
= antilog0.60
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0
2
N
8.0
0
N
N =
=
3.9
9
8.0
0
3.9
9
N = 2.00 g
Atomic Dating: not accurate for more than 50,000 yrs old due to less carbon
atoms.
➢ U -238 decomposes to Pb-206 in a series of steps
○ Half Life of U-238 is 4.5 B years
➢ After one T ½ of a sample of 1.0 g U, it would contain 0.50 g U and 0.43 g
Pb.
➢ Amount of lead is calculated from the atomic masses based on the 0.50 g U
that decomposed.
207 g/mol
Pb
X 0.50 g U
0.43 g
238 g/mol
=
Pb
U
A rock sample that contains U – 238 and Pb – 206 in the ratio
0.50 g U
is 4.5 B
0.43 g Pb
years
If the sample has larger values, it is younger. When the values are smaller, it is
older.
Nuclear Binding Energy:
E = mc2
where:
E = Energy
m = mass
e = energy
✔ Energy needed to decompose the nucleus (or the energy released when it is
formed)
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Mass Defect: Difference between the actual mass of the nucleus and the
individual sum of protons and neutrons.
Sample Problem: A 24He atom is composed of 2p, 2n and 2e. The individual
particles of the ff. masses:
p = 1.00728
n = 1.00867
e = 0.000549
Calculated Mass:
(2 x 1.007277) + (2 x 1.008665) + (2 x 0.000549) = 4.032981 amu
Mass spectrometer of a 24He atom mass = 4.002603 amu
Mass Defect
= Calculated Mass – Actual Mass
= 4.032981 – 4.002603
= 0.030378 amu (Converted to Energy)
How much energy is this?
Consider 1 mol of He atom
Total mass lost = 0.030378 g or 3.0378 x 10-5 kg
Speed of light = 2.9979 x108 m/s
E = mc2
E = (3.0378 x 10-5 kg) (2.9979 x108 m/s) 2
E = 2.7303 x 1012 kg m2/s2
In SI Units, 1 joule = 1 kg m2/s2
.:. 2.73 x 1012 J/mol
or 2.73 x 109 KJ/mol
Combustion of 1 mole CH4 liberates only 8.9 x 102 KJ/mol
Fission and Fusion Reactions
➢ Nuclear Fission: Splitting of nucleus of heavy elements into 2 or more light
elements
➢ Nuclear Fusion: formation of an element from 2 element of very low mass.
Nuclear Fission
– Extremely exothermic
– Produces about 106 times as much energy as ordinary reactions
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–
Otto Hahn and Fritz Strossman, German chemists, 1939: When U was
bombarded by slow moving neutrons, unexpected product was produced.
– Lise Meitner and Otto Frisch, discovered that it was U -235 that
absorbed neutron.
○ Fission produces chain reaction and also a source of power.
– Fissile Isotopes: capable of undergoing nuclear fission
○ Ex.: U -235, has a natural abundance of 0.72 and produces Pu -239
and U -233
– Critical Mass: the minimum amount of fissile isotopes required to sustain
the chain reaction.
○ Uranium in nuclear reactions has low purity
Nuclear Fusion
– Formation of an element of very low mass
– Requires an extremely hot temperature (40, 000, 000 0C)
– Occurs in the sun
Steps proposed to account the reaction:
12H+12H →23He+ 01n
12H+12H →13H+ 11H
13H+11H →24He+ 01n
THE PERIODIC TABLE
John Dobereiner: group elements into triads. (The Average of the atomic masses
of the first and last element will be equal to the atomic mass of the middle element)
Ex.: Fe, Co, Ni
Cl, Br, I
Basis: Occurrence of similar chemical and physical properties.
John Newlands: a scientist and a music lover; arranged based on the increasing
masses in octaves.
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Li
Be
B
C
N
O
F
Na
1
2
3
4
5
6
7
8
Do
Re
Mi
Fa
So
La
Ti
Mg
Al
Si
P
S
Cl
K
1
2
3
4
5
6
7
8
Do
Re
Mi
Fa
So
La
Ti
Do
Law of Octaves: True for lighter elements
Dmitri Mendeleev: (1834 – 1907), Russian Chemist; listed elements in increasing
atomic masses and in early 1869, reported to the Russian Chemical Society
Julius Lothar Meyer: (1830 – 1895), German Chemist; published his own
interpretation in Dec. 1869
Periodic Law: properties of elements vary periodically with their atomic number.
The Modern Periodic Table or the Long Form
2
Atomic
Number
X
Atomic Symbol
Atomic
Mass
Rows: Periods or Series (7)
Columns: Group or Family of Elements (18)
Classification of Elements
1. Representatives: Family A (Groups 1,2,13-18)
2. Transitions: Family B (3-12)
3. Inner Transition: Actinides and Lanthanides( Rare Earth Elements)
Common Names for Family of Elements
Group 1
- Alkali Metals
Group 2
- Alkaline Earth Metals
Group 17
- Halogens or Salt Formers (Greek)
Group 18
- Noble or Inert Gasses (low degree of chemical activity)
Types of Solids
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1. Amorphous Solids: arranged in a disordered form or non crystalline
structure.
Ex.: Glass and Clays
2. Crystalline Solids: particles are arranged in an orderly fashion, occupy a fix
position called crystalline lattice.
Ex.: Quartz and Table Salt
Simple Cubic
Body Centered Type
Face – centered cubic
i. Metallic
ii.Ionic
iii.Covalent Molecular
iv.Covalent Network
Elements in the Periodic Table
a. Metals
b. Non – Metals
c. Metalloids
Properties of Elements:
A. Metals
a. Shiny
b. Dense
c. Malleable
d. Ductile
e. High Melting Point
f. Good Conductor of Electricity
B. Non – Metals
a. Tend to have low densities
b. Brittle when solid
c. Most have low melting point (Some are liquid/gases at room temp)
d. Poor Conductors of electricity
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C. Metalloids
a. Exhibits the same properties of metals and non-metals
b. Electrical conductivity but not the same extent to metals; semi
conductors
METALS or METALLIC SUBSTANCES
➢ Minerals: Naturally occurring compounds of metals
○ Oxides and sulfides
○ Ore: large amount of minerals, worth mining
Physical Properties
1. High degree of electrical conductivity: delocalized electron
2. High degree of thermal conductivity: delocalized electron
3. Definite luster: presence of free electrons
4. High melting point:
Exception: Mercury (Hg) has a low melting point when solid in room temp.
Tungsten (W): has a high melting point
5. High density: distance between ions
6. Relatively strong: strong attraction between ions/ metallic bonding
7. Malleable
8. Ductile
Electron Sea Model of Metals
–
Metallic Bonding still operates after distortion.
Alloys: When different metals are melted together, stirred and cooled; mixture of
metals and solid solution
Two Types of Solid Solutions
1. Substitutional Solid Solution or Alloy
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Examples:
a. Brass: 1/3 Cu (host) replaced by Zn
b. Sterling Silver: 93% Ag; 7% Co
c. Pewter: 85% Sn, 7% Cu, 6% Bi, 2% Sb
d. Plumber’s Solder: 67% Pb, 33% Sn
1. Interstitial Solid Solution or Alloy
Examples:
a. WC or Tungsten Carbide: extremely hard, used to make cutting tools
for marking steel
b. Steel: C Atoms in the interstices of iron crystals
i. Mild Steels: 0.2% C, relatively ductile and malleable; nails, cables
and chains
ii. Medium Steels: 0.2 – 0.6% C; Relatively higher than mild steel; Rails
and structural steel beams
iii. High Carbon Steels: 0.6 – 16.5%; rough and hard; Springs, tools and
cutlery
Properties of Alloys
1. Melting point: Alloy composed of 2 metals is usually lower than the melting
point of either pure component. An extended melting range exists rather than
a sharp melting point.
Ex.: Wood’s Metal (m.p. = 70 0C)
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Bi
Pb
Sn
Cd
271.3 0C
327.4 0C
=
182 0C
320.9 0C
=
=
=
Used in sprinklers where it plug a water outlet. In case of fire,
the plug melts and water comes out to douse the fire.
1. An alloy is harder than any of the metal component
2. Electrical Conductivity: Generally lower than that of the pure metals
3. Many alloys are resistant in chemical corrosion better than their components
taken separately.

Other Alloys:
1. Stellite:
 Co and Cr;
 Used to make high speed, high cutting tools for industry
 Retain hardness even at high temperature and remains sharp
1. Stainless Steel (SS)
 Cr and Ni, Common Alloy
 Corrosion Resistant
1. 18.8 Stainless Steel
 18 % Cr ,8% Nickel,
 High corrosion resistant: Formation of Cr2O3 (Chromium (III) oxide), a
stable film. This film however is readily destroyed by Cl
 S.S. which contains Mo have higher resistance; used in salt water
1. Bronze: Cu – Zn Alloy
IONIC SUBSTANCES
✔ No discrete molecules
✔ Large Crystal Energies
✔ Hard
✔ High Melting point
✔ Brittle: repulsive charges; movement of ions
✔ Tend to shatter when struck: repulsive charges; movement of ions
✔ Solid: poor conductor of electricity
✔ Molten: good conductor of electricity
Ex.: Sodium Chloride(NaCl)
Potassium Chloride (KCl)
+
+
+
+
+
+
+
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+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
NaCl
Lattice Point: Ions
Cl- ions – Surrounded by 6 Na+
Strong Electrostatic reactions between + and – ions: Ionic Bonding
Group 1 or 2 + Group 16, 17 or 15 = Ionic Substances
COVALENT MOLECULAR SUBSTANCES
✔ Atoms sharing electrons
✔ Strong Bond
✔ Soft
✔ Low Melting point: Weak attraction
✔ Poor Conductors of Electricity: No Moving Charges
Ex.: Solid Carbon or Dry Ice
Hydrogen
O=C=O
H–H
CO2
H2
COVALENT NETWORK SUBSTANCES
✔ Held by strong covalent bonds
✔ High melting point
✔ Extremely hard
✔ Poor conductor of heat
✔ Quite brittle
✔ High boiling point
✔ Low volatility
✔ Non conductors of electricity
Ex.: Diamond, Silicon carbide or carborundum
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Allotropism: Existence of an element in more than one form, either as a result of
difference in molecular structure like O2 or O3, or as a consequence of different
packing of atom molecule in solid
Allotropic Form of Carbon:
1. Diamond: Continuous Lattice of Carbon
C
|
C
C
C
|
C
C
C
C
2. Graphite: Special Type of Covalent Network
Ex.: Lead Pencil
Properties:
a. Soft: Weak intermolecular forces
b. Conducts Electricity: Delocalized Electron
c. Possesses Metallic Luster: Delocalized Electron
1. Buckminsterfullerene or Bucky Ball
 Soot: Incomplete Combustion of Wood or Charcoal
 In 1985, Smalley and co workers, reported that large molecules have
been discovered by vaporizing molecules of graphite by lasers
 Mass Spectrometer Data: (Types of Molecules)
 Common Group of C – Magic Number
11, 15, 19 or 23
 Second group: 40- 70 C atoms
• Unusual Structure
 Arrangement of 60 – C atom was a truncated icosahedrons; a ball
shape and hollow inside
 Named after an American architect, R. Buckminster Fuller (after
geodesic domes)
 Importance: Molecule of the year 1990 – 1991
 As carrier of drugs, hollow
• Ca, La, U inserted
• Fe – cannot be inserted
 Carbon atoms can be replaced by fluorine
• Teflon: Non stick pans ; super lubricants
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LIQUID CRYSTALS: Substances that exhibit solid and liquid properties in a range
of temperature just above their melting point
 Friedrich Reintzer, Austrian botanist
○ Prepared the compound, cholesteryl benzoate
1. At room temp.: White crystalline Solid
- Heating at 145 0C, crystal structure collapses to form a turbid
liquid (liquid crystal)
- At 179 0C, turbid structure collapses to a liquid
- On cooling, the process reverses itself
2. The turbid liquid change color as temperature changes
- From red to blue with increase in temperature and vice versa
- Not all liquid crystalline compounds show color change in temp.
Kinds of Liquid Crystals
1. Nematic type
Rod like, loosely packed
2. Smectic type
3. Cholesteric
Molecule in Layer (Nematic)
Rotate spatial at a perpendicular axis
Helical structure
Uses:
Potential useful in reflecting light
– Color Mapping
Ex.: locating the vein, (warmer than the surrounding)
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