General Chemistry 1
First Quarter
Mr. Diam Salinding

Definition:
a. Anything that occupies space and has mass
b. Anything that occupies the same space and at the same time with one matters is not matter (i.e.
celestial being)

Particulate Nature of Matter
a. Every matter is made up of tiny particles
i. Atom – building block & simplest form of matter that is specific for each element and has
unique properties
ii. Molecule – combination of similar / different atoms
1. H atom + O2 molecule (diatomic)  H2O
iii. Ions – formed when atoms gain or loses an electron (i.e. Na+, Cl-)

States of Matter
Macroscopic
-
Solid
Rigid
Definite shape
Definite volume
-
Liquid
No. def. shape
Takes the shape of
container
Definite volume
-
Gas
Indefinite
volume
Indefinite shape
Microscopic




Volume (arranged in descending order): Gas, Liquid, Solid [if amount is constant]
Density (arranged in descending order): Solid, Liquid, Gas
The more space between molecules, the faster the speed of movement.
Phases Changes of Matter
Bridgit Bichara & Icee Fojas 12HA-9

Properties of Matter
a. Physical – observed without changing a substance to another substance
b. Chemical – can be observed when substance is changed to another substance
c. Intensive – amount-independent; property does not change as the amount of matter present is
changed
d. Extensive – amount-dependent; as the amount is changed, property also changes.
Examples.
1) Boiling point: physical-intensive
2) Volume: physical-extensive
3) Flammability: chemical-intensive
4) Density: physical-intensive
5) Corrosiveness: chemical-intensive
6) Color: physical-intensive
7) Mass: physical-extensive
8) Reactivity: chemical-intensive
9) Energy: physical-extensive
10) State of matter: physical-intensive
11) Absorption of light: physical-extensive
12) Photosensitivity: chemical
13) Antifreeze boils out of its radiator

Changes of Matter
a. Physical Change
i. Changes in matter that do not change the composition of a substance
ii. i.e. Sublimation of dry ice, shredding paper, evaporating alcohol, stretching a rubber band
b. Chemical Change
i. Changes that involve chemical reaction/s which results to new substances
ii. i.e. Burning wood, rusting of iron, grilling a hamburger, baking a cake, a firecracker explodes,
digesting food

Ways of Classifying Matter
a. Pure substance – a type of matter that has the same properties and the same composition
throughout a sample
b. Element – simplest pure substance and cannot be broken down any further (i.e. Na, Fe, Pb)
c. Compound – pure substance that are made up of more than one element bound together in definite
composition [constant] (i.e. H2O, NaCl, CO2)
d. Mixture – combination of different substances in varying composition
e. Homogenous mixture
i. Components cannot be distinguished from each other
ii. Appear as one substance
iii. Particles distributed evenly throughout
iv. ** Soluble component is being dissolved into a solution
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f.
v. i.e. Air, saltwater, 10k gold, alloys, a cup of coffee, mouthwash, liquid detergent
Heterogenous mixture
i. All components are visible (naked eye/microscope) because they do not mix together
ii. Phase separated
iii. Have non-uniform distribution
iv. i.e. sand & water, oil & water, pizza, vegetable soup
1. Blood – plasma is not soluble
2. Concrete – rocks are not soluble
3. Soda – gas is visible
4. Smog – suspension of smoke, fog, and dirt

Techniques in Separating Mixtures
a. Hand-picking: done manually with coarse/huge particles
b. Threshing: separate with the use of tools (i.e. separate grain from a plant)
c. Winnowing: use of wind to separate light from heavy particles (solid + solid)
d. Sieving: separate tiny from large particles (solid + solid)
e. Magnetic attraction: using a magnet to attract magnetic components
f. Sublimation: solid + gas
g. Evaporation: solid + liquid
h. Sedimentation & Decantation
i. Sedimentation: solid particles settle at the bottom
ii. Decantation: remove water using any tool or by hand to leave the particles at the bottom
i. Filtration: solid + liquid
j. Distillation: component with lower boiling point will evaporate first, enter the condenser and become
liquid in another container (liquid + liquid)
k. Centrifugation: uses centrifugal force (rapidly spins) to promote settling of particles (solid + liquid)
l. Paper Chromatography: separate colored substances (i.e. determine components in ink)

English System vs Metric System
Measurement
Length
Mass
Volume
Temperature
Time

English
Yard / Inch
Ounce / Pound
Quart
Fahrenheit
Second
System
Meter / Centimeter
Gram / Kilogram
Liter
Celsius / Kelvin
Second
SI Units (International System)
7 Fundamental Units
Physical Quantity (Dimension) Unit Name
Mass
Kilogram
Length
Meter
Time
Second
Temperature
Kelvin
Electric current
Ampere
Amount of substance
Mole
Luminous intensity
candela
Unit Abbreviation
kg
m
s
K
A
mol
cd
a. Conversion units
12 inches
3 feet
5, 280 feet
16 ounces
4 quarts
1 foot
1 yard
1 mile
1 pound
1 gallon
Bridgit Bichara & Icee Fojas 12HA-9
b. The SI System is based on the number 10
c. Common Prefixes used with SI units
Prefix
Abbreviation Meaning
Tera
T
1012
Giga
G
109
Mega
M
106
Kilo
k
103
Hecto
h
102
Deca
da
101
Deci
d
10−1
Centi
c
10−2
Milli
m
10−3
Micro
µ𝑎
10−6
Nano
n
10−9
Pico
p
10−12
Femto
f
10−15
Example
1 terameter (Tm) = 1 𝑥 1012 𝑚
1 gigameter (Gm) = 1 𝑥 109 𝑚
1 megameter (Mm) = 1 𝑥 106 𝑚
1 kilometer (km) = 1 𝑥 103 𝑚
1 hectometer (hm) = 1 𝑥 102 𝑚
1 decameter (dam) = 1 𝑥 101 𝑚
1 decimeter (dm) = 0.1 𝑚
1 centimeter (cm) = 0.01 𝑚
1 millimeter (mm) = 0.001 𝑚
1 micrometer (µm) = 1 𝑥 10−6 𝑚
1 nanometer (nm) = 1 𝑥 10−9 𝑚
1 picometer (pm) = 1 𝑥 10−12 𝑚
1 femtometer (fm) = 1 𝑥 10−15 𝑚
d. Temperature
i. Kelvin Scale
1. Lowest temperature possible (absolute zero) is zero Kelvin.
2. Absolute zero: 0K = -273.15°C.
ii. Celsius Scale
1. Water freezes at 0°C and boils at 100°C
2. To convert: K = °C + 273.15
3. 25°C  standard room temperature
4. 0°C  standard temperature
iii. Fahrenheit Scale
1. Water freezes at 32°F and boils at 212°F
2. To convert:
5
9
°C = (°F − 32) 𝑜𝑟 °F = (°C) + 32
9
5
Example. Convert the following temperature to °C
77°K → 77°K − 273.15 = −196.15°C
5
4.2°F → (4.2 − 32) = −15.44°C
9
−96°K → −96°K − 273.15 = −369.15°C
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e. Derived units
i. Obtained by multiplication or division of one or more of the 7 SI units.
1. Example 1: Speed
𝑢𝑛𝑖𝑡𝑠 𝑜𝑓 𝑑𝑖𝑠𝑡𝑎𝑛𝑐𝑒
𝑈𝑛𝑖𝑡𝑠 𝑜𝑓 𝑣𝑒𝑙𝑜𝑐𝑖𝑡𝑦 =
𝑢𝑛𝑖𝑡𝑠 𝑜𝑓 𝑡𝑖𝑚𝑒
𝑚𝑒𝑡𝑒𝑟𝑠
=
𝑠𝑒𝑐𝑜𝑛𝑑𝑠
= 𝑚/𝑠
2. Example 2: Volume
a. The units for volume are given by the (units of length)3
i. SI unit for volume is 1 𝑚3
b. We usually use 1 mL – 1 𝑐𝑚3
c. Other volume units:
i. 1 𝐿 = 1 𝑑𝑚3 = 1000 𝑐𝑚3 = 1000 𝑚𝐿
3. Example 3: Density
a. Used to characterize substances
b. Defined as mass divided by volume:
𝑚𝑎𝑠𝑠
𝐷𝑒𝑛𝑠𝑖𝑡𝑦 =
𝑣𝑜𝑙𝑢𝑚𝑒
c. Units: 𝑘𝑔/𝑚3
d. Originally based on mass (the density was defined as the mass of 1.00 g of
pure water).
Densities of Some Selected Substances at 25°𝐂
Substance
Density (𝒈/𝒄𝒎𝟑 )
Air
0.00
Balsa wood
0.16
Ethanol
0.79
Water
1.00
Ethylene glycol
1.09
Table sugar
1.59
Table salt
2.16
Iron
7.9
Gold
19.32

Uncertainty in Measurements
a. Exact number – values are known to be exact
b. Inexact number – values having some uncertainty
Exact
Inexact
1000 g/kg
Ruler measure
2.54 cm/in
Temperature reading
12/dozen
Volume or Mass
Any conversion factor
Etc.
c. All scientific measures are subjected to error (i.e. equipment errors, human errors)
d. These errors are reflected in the number of figures reported for the measurement.
e. Different measuring devices have different uses and different degrees of accuracy/precision.
f. The number of digits reported in a measurement reflect the accuracy of the measurement and the
precision of the measuring device.
g. Precision – measure of how closely individual measurements agree with one another
h. Accuracy – how closely individual measurements agree with the correct, or “true,” value

Significant Figures
a. Refer to digits that were measured
b. All digits of a measured quantity, including the uncertain one, are called significant figures
c. When rounding calculated numbers, we pay attention to significant figures so we do not overstate
the precision of our answers
d. Rules:
i. All nonzero digits are significant.
423.444
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ii. Zeroes between two significant figures are themselves significant.
42300045
42,340.0025
iii. Zeroes at the beginning of a number are never significant.
00042345.0
0.00048
iv. Zeroes at the end of a number are significant if the number contains a decimal point.
423,000 versus: 423,000. or 423,000.000

Rules for rounding off
a. If the digit removed is 5 or more than 5, the preceding number is increased by 1.
b. If the digit removed is less than 5, the preceding number is not changed.
i.e.
1.496  1.50
1.494  1.49

Significant Figures in Calculations: Addition, Subtraction, Multiplication and Division
a. For addition and subtraction, the result has the same number of decimal places as the measurement
with the fewest decimal places.
b. For multiplication and division, the result contains the same number of significant figures as the
measurement with the fewest significant figures.
i.e.
1.) 46.32
9.1
+ 12.316
67.736 → 67.7
2.) 102 − 26.73 = 75
3.2 𝑥 5.01
= 3.4
4.695
3
3.24 𝑥 10
2
4
4. )
−2 = 2.58 𝑥 10 = 6.37 𝑥 10
5.11 𝑥 10
3. )
5. ) [3.46 𝑥 107 − 6.25 𝑥 105 ] + (467)(32) = 3.4 𝑥 107

Dimensional Analysis
a. Involves the correct use of conversion factors to change one unit into another.
b. Conversion factor = a fraction whose numerator and denominator are the same quantity expressed
in different units
c. When we multiply a quantity by a conversion factor, the units multiply and divide as follows:

Conversion Factors
Length:
1 inch (in) = 2.54 centimeters (cm)
1 foot (ft) = 12 inches (in)
1 mile (mi) = 1.60934 kilometers (km)
1 meter (m) = 1.09361 yards (yd)
1 km = 1000 m
1 m = 100 cm
1 dm = 10 cm
Mass:
1 kilogram (kg) = 2.20462 pounds (lbs)
1 pound (lb) = 16 ounces (oz)
1 metric ton = 1000 kilograms
Volume:
1 mL = 1 cm3
1 L = 1 dm3
1 US gallon (gal) = 128 US fluid ounce (fl. oz.)
1 US gallon = 3.78541 Liters
1 UK gallon = 4.54609 Litres
1 Liter = 1.05669 quarts
Time:
1 year = 365 days
1 day = 24 hours
1 hour = 60 minutes
1 minute = 60 seconds
Bridgit Bichara & Icee Fojas 12HA-9

Democritus – described the material world as made up of tiny indivisible particles that they called
atomos, meaning “indivisible” or “uncuttable”

Plato and Aristotle
a. Said there can be no ultimately indivisible particles
b. Materials are infinitely divisible
c. Most accepted this claim during their time

Antoine Lavoisier
a. Discovered 19 metals and 7 nonmetals
b. Father of Modern Chemistry
c. Responsible for Law of Conservation of mass  cannot be created nor destroyed

Amadeo Avogadro – responsible for Law of Combining Volume

Joseph Proust
a. Responsible for Law of Constant Composition
b. Defined compound

John Dalton’s Atomic Theory of Matter
a. Each element is composed of extremely small indivisible particles called atoms.
b. All atoms of a given element are identical but the atoms of one element are different from the atoms
of all other elements.
i. i.e. Oxygen is made up of oxygen atoms only which all have the similar properties
ii. i.e. Hydrogen is made up of hydrogen atoms only which all have the similar properties
iii. But the atoms that compose oxygen are different in physical and chemical properties from
the atoms that composed hydrogen.
c. Atoms of one element cannot be changed into atoms of a different element by chemical reactions;
atoms are neither destroyed nor created (Law of Conservation of Mass)
i. i.e. 1g of 2 H2O + 1g of O2  2g of 2 H2O
d. Compounds are formed when atoms of more than one element combine; a given compound has the
same relative number and kind of atoms (Law of Constant Composition)
i. Same number of atoms in products and in reactants
ii. i.e. 2 H2O + O2  2 H2O

Key concept: Particles with the same charge repel, whereas opposite charges attract


This experiment gave way to the discovery of the electrons by measuring the charge to mass ratio.
Bridgit Bichara & Icee Fojas 12HA-9

Process
a. It was a vacuum environment (Cathode to Anode tube)
b. H gas was placed in the environment.
c. Electricity was introduced (there was a potential difference)  light was formed
d. A phosphor material (orange circle in the figure) was added and a hole was created in the anode
e. They charged it again and the first ray largely deflected from the negative. If the ray was deflected
from the negative, following the like charges repel, then the particles were negatively-charged 
called the particles corpuscles
f. From this first ray, most of the electrons were already lost.
g. When they charged it again, the ray was slightly deflected from the positive. Ergo there were also
positively charged particles which they knew came from the H+ ions

Charge to Mass ratio
a. Deflection is directly proportional to charge and indirectly proportional to mass
∆𝑥 ∝
𝑒−
𝑚−
𝑤ℎ𝑒𝑟𝑒 ∆𝑥 = deflection
∆𝑥 ∝
𝑒𝐻+
𝑚𝐻+
b. The deflection of the negatively-charged particles was greater than that of the positive.
∆𝑥− > > > ∆𝑥𝐻+
c. The mass of H+ ion is greater than that of the negatively charged particle.
𝑒−
∆𝑥−
𝑀𝐻+
𝑚−
= 𝑒
=
𝐻+
∆𝑥𝐻+
𝑀−
𝑀𝐻+
𝑀𝐻+ > > > 𝑀−
d. Ergo, the lighter the element, the faster and bigger the deflection.
e. At the time, finding a lighter element than Hydrogen, which was the presumed lightest element, was
a big discovery. It was in contrast to Dalton’s theory  there is a smaller particle than an atom.

Introduced the Plum-Pudding Model
a. Positively-charged pudding with negatively-charged plums
Bridgit Bichara & Icee Fojas 12HA-9


This is to measure the actual magnitude/charge of electrons





Process: It was subjected to x-ray to ionize the oil  negatively charged oil
Ef = Gf  determined by mass
The higher the voltage, the higher the repulsion = oil rises up toward the positive plate
The lower the voltage, the lower the repulsion = oil sinks down toward the negative plate
Goal: Adjust to the right amount of voltage for the particle to suspend in the middle
Zepto Columns (10-21)
160 (1 x 160)
320 (2 x 160)
480 (3 x 160)
640 (4 x 160)
800 (5 x 160)
a. They saw an arithmetic pattern in the zepto columns  a difference of 160 per column

Conclusion: Charge of an electron is -160 x 10-21 or -1.60 x 10-19



Materials:
a. RaBr2 – radioactive source which emits alpha particles which are heavy and charged
b. Geiger Counter – detect emissions of radiation through ticking sounds
To understand alpha particles, they got a thin gold foil (thinner than hair) and tried if gold can block the
alpha particles. Yet, before and after the use of the gold foil, there were still 132,000 particles per minute
 so they approved JJ Thomson’s statement that electron particles pass through
Bridgit Bichara & Icee Fojas 12HA-9

Ernest Marsden, student of Rutherford, was curious so he placed the Geiger counter at a different angle.
a. Without gold foil, no ticking sound.
b. With gold foil, with ticking sound  20 particles per minute
c. Disproves JJ Thomson’s plum pudding model.

Conclusion
a. Nucleus caused alpha particles to bounce back (Discovery of Nucleus)
b. If the alpha particles are positive, following the like charges repel, there must be positively charged
particles inside the nucleus.  discovery of protons
c. The alpha particles turned out to be He++


Diameter of atom: 10-10 m
Diameter of nucleus: 10-14 m



Used polonium (radioactive substance) and Beryllium
Used gamma particles that showed to have no charge  neutrons
Particle
Proton
Neutron
Electron
Charge
Positive (1+)
None (neutral)
Negative (1-)
Mass (amu)
1.0073
1.0082
omoooo
e<p<n
a. 1 amu – 1.66054 x 10-24 g
b. Atoms diameter – 1-5 Angstrom (Å)
c. 1Å = 1 x 10-10 m

Atomic number – the number of protons/electrons in an atom

Mass number
a. Number of protons + number of neutrons
b. Atomic number + number of neutrons
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
Isotopes
a. Same atomic number, different atomic mass
b. Same no. of protons/electrons, different no. of neutrons
c. Protium, Deuterium, Tritium have the same number of protons but differ in atomic mass.

Atomic weight – is the element’s average atomic mass
Atomic weight = ∑[(isotope mass) x (fractional isotope abundance)]
overall isotopes of the element
a. i.e. Naturally occurring carbon, for example, is composed of 98.93% 12C and 1.0736 13C. The
masses of these isotopes are 12 amu (exactly) and 13.00335 amu, respectively, making the atomic
weight of carbon
(0.9893)(12 amu) + (0.0107)(13.00335 amu) = 12.01 amu

Periodic Table
a. Arrangement of elements in order of increasing atomic number, with elements having similar
properties placed in vertical columns, is known as the periodic table.
b. Many elements show strong similarities to one another.
i. i.e. Lithium (Li), sodium (Na), and potassium (K)  soft & very reactive metals
ii. i.e. Helium (He), neon (Ne), and argon (Ar)  very nonreactive gases.
c. Elements in a group often exhibit similarities in physical and chemical properties
i. Group 18: “coinage metals” – copper (Cu), silver (Ag), and gold (Au). These elements are
less reactive than most metals, which is why they have been traditionally used throughout
the world to make coins.
d. Period – horizontal rows of periodic table
e. Groups – vertical columns

Molecular compounds
a. Occurs between nonmetals (bond between them is covalent)
b. Compounds composed of molecules containing more than one type of atom
c. i.e. Methane (CH4), Water (H2O), Table sugar (C12H22O11), Pure ethyl alcohol (CH3CH2OH)

Molecular Formulas – chemical formulas that indicate the actual numbers of atoms in a molecule
a. i.e. H2O, NH3, CH4,H2O2, C6H12O6

Empirical Formulas – chemical formulas that give only the relative number of atoms of each type in a
molecule  simplify and reduce the subscripts
a. H2O2  HO
b. C6H12O6  CH2O
c. NH3  NH3

Ions – formed if electrons are removed from or added to an atom  loss or gain
a. Cation – positively charged ion  metals: lose electrons: positive
b. Anion – negatively charged ion  nonmetals: gain electrons: negative
c. How many protons, neutrons, and electrons are in the ff:
2−
i. 34
16𝑆
3+
ii. 41
21𝑆𝑐

Ionic compound – metal + nonmetal (i.e. NaCl, K2SO4, RaBr2)
Bridgit Bichara & Icee Fojas 12HA-9

Predicting Ionic Charges
a. Many atoms gain or lose electrons to end up with the same number of electrons as the noble gas
closest to them in the periodic table.
b. We might deduce that atoms tend to acquire the electron arrangements of the noble gases because
these electron arrangements are very stable. Nearby elements can obtain these same stable
arrangements by losing or gaining electrons.
i.
ii.
iii.
iv.
+1: one step back to acquire arrangement of noble gas
+2: two steps back to acquire arrangement of noble gas
-1: one step forward to acquire arrangement of noble gas (needs to gain electrons)
As you go to the right of the periodic, it gets more negative.
Binary Compounds = 2 elements
1. Metal + Metal: Fixed oxidation number / charge [Group IA / IIA]
a. Rule
i. Name the 1st element by its element name
ii. Name the 2nd element with the suffix changed to -ide
iii. If subscripts are reducible, then reduce.
b. Examples
i. K2O --- potassium oxide
ii. Ba3N2 --- barium nitride
iii. Al2O3 --- aluminum oxide
iv. Ca2O2 --- CaO --- calcium oxide
2. Metal + Metal: Variable oxidation number / charge
a. Rule
i. Stock System – use of roman numerals to indicate the O.N / charge of the metal
ii. Classical System
1. Lower ON / charge  -ous
2. Higher ON / charge  -ic
3. i.e. Cuprum (Cu), Stannum (Sn), Plumbum (Pb), Ferrum (Fe)
b. Examples
Stock
Classical
SnCl4 tin (IV) chloride
stannic chloride
FeBr3 iron (III) bromide ferric bromide
Co2S3 cobalt (III) sulfide cobaltic sulfide
SnCl2 tin (II) chloride
stannous chloride
FeBr2 iron (II) bromide
ferrous bromide
CoS
cobalt (II) sulfide cobaltous sulfide
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3. Nonmetal + Nonmetal
a. Rule
i. Name the 1st element by its element name
ii. Name the 2nd element by changing the suffix to –ide
iii. Use Greek prefixes to indicate the number of the 1st and 2nd element
12345-
mono
di
tri
tetra
penta
678910 -
hexa
hepta
octa
nona
deca
iv. The prefix mono- for the 1st element is omitted
v. If the name of the 2nd element starts with a vowel, the a or e in the prefix will be dropped.
b. Example
i. N2O5 --- dinitrogen pentoxide
Ternary Compounds
1. Metal + Polyatomic Anion: Fixed oxidation number / charge [Group IA / IIA]
a. Examples
sodium sulfate
Na+1 + SO4-2
Na2SO4
barium nitrate
Ba+1 +
NO3-1
Ba(NO3)
sodium sulfite
Na+1 + SO3-2
Na2SO3
barium nitrite
Ba+1 + NO2-1
Ba(NO2)2
2. Metal + Polyatomic Anion: Variable oxidation number / charge
a. Examples
Cobaltous phosphate
Cobalt (II) phosphate
Co+2 + PO4-3
Stannous phosphate
Tin (II) phosphate
Sn+2 +
PO4-3
Pb+4 +
SO4-2
Co3(PO4)2
Sn3(PO4)2
Plumbic phosphate
Lead (IV) sulfate
Pb4(SO4)2  Pb(SO4)2
b. For mercury (I), you cannot reduce the subscript no matter the case.
i. i.e. mercury (I) acetate --- Hg2 (C2H3O2)2  do not reduce!
Bridgit Bichara & Icee Fojas 12HA-9
Inorganic Acids
1. H+ + Nonmetal
a. Rule
i. Dry Form – same as naming binary compounds
ii. Aqueous Form – hydro + element + -ic + acid
b. Example # 1
i. Dry Form: H2S --- hydrogen sulfide
ii. Aqueous Form: H2S (aq.) --- hydrosulfuric acid
c. Example # 2
i. Dry Form: HI --- hydrogen iodide
ii. Aqueous Form: HI (aq.) --- hydroiodic acid
2. H+ + Polyatomic Anion: Variable oxidation number / charge
a. Rule
Dry Form
-ite
-ate
Aqueous Form
-ous
-ic
b. Example # 1
i. Dry Form: H2SO4 --- hydrogen sulfate
ii. Aqueous Form: H2SO4 (aq.) --- sulfuric acid
c. Example # 2
i. Dry Form: H2SO3 --- hydrogen sulfite
ii. Aqueous Form: H2SO3 (aq.) --- sulfurous acid
d. Example # 3
i. Dry Form: H3PO4 --- hydrogen phosphate
ii. Aqueous Form: H3PO4 (aq.) --- phosphoric acid
Hydrates (w/ H2O)
1. Examples
a. Ca(NO3)2 ∙ 2H2O --- calcium nitrate dehydrate / 2-hydrate
b. Cu2SO4 ∙ 5H2O --i. cuprous sulfate pentahydrate / 5-hydrate
ii. copper (I) sulfate pentahydrate / 5-hydrate
Bridgit Bichara & Icee Fojas 12HA-9

For oxygenated ions:
a. Hypo + -ite = O1
b. -ite = O2
c. -ate = O3
d. Per + -ate = O4

Rutherford’s Nuclear Model of Atom
a. Shows that an atom is mostly empty space containing three
sub-atomic particles called protons, neutrons and electrons.
i. The protons and neutrons are found in the nucleus at
the center of the atom. The nucleus is very much
smaller than the atom as a whole.
ii. The electrons are found outside the nucleus.

Bohr’s Atomic Model
a. Electrons orbit the nucleus like planets orbit the sun
Bridgit Bichara & Icee Fojas 12HA-9

Quantum Mechanical Model
a. Most advanced and accurate model of the atom, used today by
chemists and physicists
b. In this model, electrons do not exist as tiny points inside the atom,
but instead surround the nucleus in a form resembling a cloud

Heisenberg’s Uncertainty Principles
a. The quantum mechanical model of the atom uses complex shapes
of orbitals (sometimes called electron clouds), volumes of space in
which there is likely to be an electron. So, this model based on probability rather than certainty 
quantum numbers

Definition:
a. An orbital is described by a set of four quantum numbers
b. The 4 quantum numbers specify the energy and location of electrons around a nucleus.

Principal quantum number (n)
a. Describes the energy level/shell on which the orbital resides
b. The values of n are integers ≥1
c. As n becomes larger, the electron is further from the nucleus.
d. The closer the nucleus, the lower the energy.

Azimuthal or angular momentum quantum number (l)
a. Defines the shape of the orbitl.
b. Allowed values of l are integers ranging from 0 to n – 1
c. The atom’s energy level/shells contains sublevels/subshells which designate the orbital shape. Each
subshell has a letter designation:
Value of l
Type of orbital

0
s
1
p
2
d
3
f
Magnetic quantum number(ml)
a. Describes the three-dimensional orientation of the orbital.
b. Values are integers ranging from –l to l:
-l ≤ ml ≤ l
Bridgit Bichara & Icee Fojas 12HA-9
c. Therefore, on any given energy level, there can be up to 1s orbital, 3p orbitals, 5d orbitals, 7f
orbitals, etc.
d. Orbitals with the same value of n form a shell
e. Different orbital types within a shell are subshells
f. Total number of orbital is the total number of subshell (total number of orbitals = n2)

Spin quantum number (ms)
a. Two electrons in the same orbital do not have exactly the same energy
b. The “spin” of an electron describes its magnetic field which affects its energy
1
1
c. The spin quantum number has only 2 allowed values: + 2 , − 2
i. If the last electron has a spin going up, then it is positive.
ii. If the last electron has a spin going down, then it is negative.

Summary
Name
Principal
Angular Momentum

Symbol
n
l
Magnetic
ml
Spin
ms
Permitted Values
Positive integers (1,2,3…)
Integers from 0 to n – 1
Integers from –l to 0 to +l
1 1
+ ,−
2 2
Property
Orbital energy (size)
Orbital shape (0,1,2, and 3
correspond to s,p,d and f
orbitals, respectively)
Orbital orientation in space
Direction of e- spin
Sample Questions
a. Which one of the following is not a valid value for the magnetic quantum number of an electron in a
5d subshell? Answer: 3
i. 5d subshell: n=5; l=2 (from d); ml = -2, -1, 0, 1, 2
b. An electron cannot have the quantum numbers n = __, l = __, ml = __ Answer: 1, 1, 1
c. Which set of three quantum numbers (n,l,ml) corresponds to a 3d orbital? Answer: A
Bridgit Bichara & Icee Fojas 12HA-9
i. 3d subshell: n=3; l=2; ml = -2, -1, 0, 1, 2
d. Which one of the following represents an acceptable set of quantum numbers for an electron in an
atom? (n, l, ml, ms) Answer: B
1. 2, 2, -1, - ½  does not satisfy that l must be n – 1
2. 1, 0, 0, ½
3. 3, 3, 3, ½  does not satisfy that l must be n – 1
4. 5, 4, -5, ½  does not satisfy that ml must be ≥ n
5. 3, 3, 3, - ½  does not satisfy that l must be n – 1

An electron configuration of an atom is a particular distribution of electrons among available sub
shells
a. The notation for a configuration lists the sub-shell symbols sequentially with a superscript indicating
the number of electrons occupying that sub shell
b. i.e. Lithium (atomic number 3), has two electrons in the “1s” sub shell and one electron in the “2s”
sub shell 1s2 2s1.
2 Ways to write electron configurations
spdf notation
Orbital box notation

Rules for filling orbitals:
a. Bottom-up
i. Aufbau’s Principle
1. Electrons occupy the lowest energy levels available
2. A scheme used to reproduce the ground state
electron configurations of atoms by following the
“building up” order.
ii. Order in which all the possible sub-shells fill with electrons
1s, 2s, 2p, 3s, 3p,
(Si Shenly Pumasok sa Pinto)
4s, 3d, 4p, 5s, 4d,
(Si Dalida Pumasok sa Door)
5p, 6s, 4f, 5d,
(Papano si Fafa Diam)
6p, 7s, 5f, 5d, 7p
(Papano si Fafa Diam, Papano)
b. Fill orbitals singly before doubling up
i. Hund’s Rule: lowest energy arrangement of electrons in a
subshell is obtained by putting electrons into separate orbitals of the subshell with the same
spin before pairing electrons
ii. Ex. Oxygen
1s
2s
2p
c. Paired electrons have opposite spin (Pauli’s exclusion principle)
i. Pauli’s exclusion principle: An orbital can hold at most two electrons, and then only if the
electrons have opposite spins.
ii. Hence, no two electrons can have the same four quantum numbers.
iii. The maximum number of electrons and their orbital diagrams are:
Bridgit Bichara & Icee Fojas 12HA-9
Subshell

Number of
Orbitals
Maximum Number
of Electrons
s (l = 0)
1
2
p (l = 1)
3
6
d (l = 2)
5
10
f (l = 3)
7
14
Condensed Electronic Configuration / Shorthand Notation Practice
a. [Noble Gas core] = higher energy electrons
b. i.e.
i. Aluminum: 1s22s22p63s23p1  [Ne] 3s23p1]
ii. Calcium: 1s22s22p63s23p64s2  [Ar] 4s2
iii. Nickel: 1s22s22p63s23p64s23d8  [Ar] 4s23d8
iv. Iodine: [Kr] 5s24d105p5
v. Astatine: [Xe] 6s24f145d106p5
c. There are a few exceptions to the building-up order prediction for the ground state.
i. Chromium (Z = 24) and Copper (Z=29) have been found by experiment to have the following
ground-state electron configurations:
1. Cr: 1s22s22p63s23p64s13d5
2. Cu: 1s22s22p63s23p64s13d10
i. Some irregularities occur when there are enough electrons to half-fill s and d orbitals on a
given row  this occurs because the s and d orbitals are very close in energy
Bridgit Bichara & Icee Fojas 12HA-9
d. Exercise: Give the set of quantum numbers for the last electron occupying the last subshell of:
i. Vanadium: n = 3, l = 2, ml = 0, ms = + ½
ii. Calcium: n = 4, l = 0, ml = 0, ms = - ½

Unpaired electrons in orbitals gives rise to paramagnetism and is attracted to a magnetic field.
Diamagnetic species contain all paired electrons and is “repelled” by the magnetic field.
a. Diamagnetic atoms or ions:
i. All e- are paired
ii. Weakly repelled in a magnetic field
b. Paramagnetic atoms or ions:
i. Unpaired e- exist in an orbital
ii. Attracted to an external magnetic field.
c. Metals loose electrons (oxidized) and become cations. Nonmetals gain electrons (reduced) and
become anions. The electronic configuration of each reflects this change in the number of electrons.
i. Metals lose electrons so that cation has a noble-gas outer electron configuration.
ii. Non-metals gain electrons so that anion has a noble-gas outer electron configurations.
iii. Metals and non-metal ions tend to form electronic states closest to their nearest noble gas
configuration.

Isoelectronic species are two different elements (usually with charged electrons) with the same
electronic configuration – but not the same nuclear configuration.
a. When a transition-metal cation is formed from an atom of a transition metal, electrons are remove
first from the ns orbital, then from the (n-1) orbital.
Bridgit Bichara & Icee Fojas 12HA-9

The periodic law states that when the elements are arranged by atomic number, their physical and
chemical properties vary periodically

Atomic Radius
a. Two factors determine the size of an atom
i. Principal quantum number (n)
1. The larger the n, the larger the size of the orbital
ii. Effective nuclear charge (Zeff):
1. Positive charge an electron experiences from the nucleus minus any “shielding
effects” from intervening electrons.
b. Electrons in elements are categorized either as:
i. Inner core electrons
1. electrons residing in the lower n shells of an element
2. located closer to the nucleus
3. “shield” or “screen” outer electrons from the positive charge of the nucleus
ii. Outer core or Valence e1. Total number of e- in the highest n-value shell
2. No. of Valence e- = Group Number
3. Screening impacts:
a. Alters the energy levels spacing and ordering in many electron atoms
b. Outer e- screened by inner electrons
c. Zeff is the electrostatic force (force of attraction between an electron and the nucleus) left by the
outer valence electrons takin into “shielding” by internal core electrons.
To a good approximation: effective nuclear charge Zeff is given by:
Zeff = Z – core eWhere:
Zeff = Effective nuclear charge
Z = Number of protons in atom
Core e- = Number of inner core non-valence electrons
Larger Zeff means more “pull” or electrostatic force between
nucleus and electrons
d. Atomic Radii decrease across a period because the effective nuclear increases.
i. The closer the outer valence electrons, the stronger the Zeff
ii. The closer the outer valence electrons, the smaller the atomic radii.
Bridgit Bichara & Icee Fojas 12HA-9
e. Exercise: Rank each set of main group elements in order of decreasing atomic size:
i. Ca, Mg, Sr  Sr > Ca > Mg
ii. K, Ga, Ca  K > Ca > Ga
iii. Br, Rb, Kr  Rb > Br > Kr
iv. Sr, Ca, Rb  Rb > Sr > Ca
f. For charged electrons:
i. Anions  bigger than their ground state atom  gain electrons  negative
ii. Cations  smaller than their ground state atom  lose electrons  positive
iii. Greater cation charge  smaller and vice versa
g. Exercise: Rank Ions by size.
i. Ca2+, Sr2+, Mg2+ , Sr2+ > Ca2+ > Mg2+
ii. K+, S2-, Cl-  S2- > Cl- > K+
iii. Au+, Au3+  Au+ > Au3+
h. Tip: Follow charge if different charges, follow trend if same charges.

Ionization Energy
a. The minimum energy (kJ/mol) required to remove 1 mole of e- from one mole of a gaseous atom in
its ground state.
b. The increasing effective nuclear charge (Zeff) and its impact on atomic radius can help us understand
the trend in ionization energies of elements.
c. The ability to lose e-: To lose electrons, there is a need to bombard the electron charge bond
between electron and nucleus  energy to destroy the energy between electron & nucleus
d. We can remove more than 1 electron from a ground state atom. It requires more energy to remove
subsequent electrons.
i. l1 + X(g)  X+(g) + el1 first ionization energy
ii. l2 + X(g)  X2+(g) + el2 first ionization energy
iii. l3 + X(g)  X3+(g) + el3 first ionization energy
l1 < l2 < l3
Bridgit Bichara & Icee Fojas 12HA-9
e. Exercise: Rank the elements in each set in order of decreasing lE1:
i. Kr, He, Ar

He > Ar > Kr
ii. Sb, Te, Sn

Te > Sb > Sn
iii. K, Ca, Rb

Ca > K > Rb
iv. I, Xe, Cs

Xe > I > C

Metallic behavior increases as we move down a group and from left to right on the periodic table.
a. Metals have low ionization energy
b. Tend to be oxidized to metal ions.
c. Charge is group numbers

Electron affinity
a. Is the energy required to add (reduce) 1 mole of e- to an atom in the gas state to form an anion. It’s
a measure of an atom’s ability to “accept” an eb. The ability to accept ec. Largest (most negative) for Chlorine and Fluorine (i.e. like to gain electrons = reduced).
d. Ionization energy and electron affinity are almost the same therefore they follow the same trend. It is
just that electron affinity is more negative because it accepts ee. If there is a decrease in size of atomic radii, more energy is needed.

Electronegativity
a. Define as the ability of an atom in a molecule to attract electrons to itself
Bridgit Bichara & Icee Fojas 12HA-9