AP Chemistry

Chapter 2
Atoms, Molecules, and Ions
 UNIT 2 LEARNING TARGETS
2.1) I can discuss atomic structure (both historically and currently) and how it relates to atomic
properties.
a.I can describe the basic postulates of Dalton’s Atomic Theory.
b.I can describe the key historical discoveries of J.J. Thomson, Millikan, Curie, Rutherford, and Chadwick
that led to the discovery of the structure of the atom.
a.I can describe the structure of the atom in terms of protons, neutrons, and electrons.
b.I can list the relative charge, mass, and location of protons, neutrons, and electrons.
c.I can use isotope notation to express the subatomic composition of atoms.
d.I can calculate average atomic mass from natural abundance of isotopes and masses of individual atoms.
e.I can explain how ions are formed and determine the number of electrons gained or lost based on ionic
charge.

2.2) I can use the periodic table as a reference tool.

a.I can describe how elements are organized in the periodic table leading to periods and groups.
b.I can identify the location of metals, nonmetals, and metalloids.
c.I can use the periodic table to predict charges of ions.

2.3) I can write names and formulas for simple, inorganic compounds.
a.I can identify, name , write formulas for ionic compounds.
b.I can identify, name, and write formulas for molecular compounds.
c.I can identify, name, and write formulas for acids.
 Atoms
● Building blocks of matter
● Smallest unit of an
element that retains the
element’s properties
● Atoms are made up of
still smaller particles
● Concept/model of the
atom has evolved over Individual carbon atoms in graphite
imaged by STM
2500 years.
 DEMOCRITUS
Philosophical concept of atoms
- Greek philosopher Leucippus and student Democritus
- all matter is made up of tiny, indivisible particles, or atoms,
- 5th century B.C.
- atom comes from Greek word atomos, which means “indivisible.”
- too small to be seen, unchangeable, and indestructible,
- completely solid, with no internal structure,
- variety of shapes and sizes: different kinds of matter.
- Color, taste, and other qualities were thought to be “atomic.”

Aristotle argued against the existence of such atoms

-his world view “won out”
-atomic philosophy was largely dismissed for centuries
 ARISTOTLE’S THEORY OF MATTER
● All substances combinations of elements and elemental qualities.
● The elements are: fire, water, earth, and air.
● (Aristotle added later another "element" - Ether which was a perfect
substance and what the heavenly bodies are composed of)
● The qualities are: hot, cold, wet, dry.
● The qualities define the character of "elements".
● Fire was seen as ideal mixture of hotness & dryness.
● One element could be changed into another like mixing solutions.
● Combustion and other chemical reactions were considered a type of
motion.
 ANTOINE LAVOISIER
● Element concept provided the foundation for Dalton's
atomism.
○ “Elements are the substances we have not discovered means for
separating. They are all the substances into which we are
capable to reduce [in weight] bodies by decomposition.”

→ Defined elements as substances that could not be

broken into simpler substances through chemical means.

● Law of Conservation of Mass

○ “Nothing is created, nothing is lost, everything is transformed.”
○ Basis of quantitative chemistry
○ Weight becomes the tool for determining if a product is simpler or
more complex than an ingredient (a reactant).
JOHN DALTON
John Dalton (1766-1844):
• One of the earlier scientists to
work on the understanding of
matter.
• At the hub or focus of his
assumption is that matter is
discontinuous.
• The four assumptions that he
made make up his theory of
atoms.
 DALTON’S ATOMIC THEORY
New System of Chemical Philosophy 1808

1) Matter is composed of small indivisible particles called

atoms.
● The atom is the smallest unit of an element that enters
into chemical combination.
2) An element is composed entirely of one type of atom.
● The properties of all the atoms of one element are
identical (particularly the mass) and are different from
those of any other element.
3) Compounds are formed by a combination of atoms of
two or more different elements joined together.
4) Chemical reactions simply change the way the atoms
are joined together. Atoms do not change their identities
in chemical reactions.
 EXPERIMENTAL EVIDENCE FOR DALTON’S ATOMIC THEORY
● Different elements were made up of different atoms and this
explained the different properties of the elements.

● Law of Conservation of Mass (Lavoisier 1785) was consistent

with “unbreakable” atoms that underwent rearrangements in
chemical reactions.
● Law of Definite Proportions (Proust 1799)
○ all compounds have definite ratio by mass of elements of which they are
composed
○ nature has a definite “recipe” for every compound
 EXPERIMENTAL EVIDENCE FOR DALTON’S ATOMIC THEORY
● Law of Multiple Proportions (Dalton 1808)
○ If two elements form more than one compound together, then the ratios of the
masses of the second element which combine with a fixed mass of the first
element will be ratios of small whole numbers.
○ oxygen and carbon combined to make two compounds
○ each has its own particular mass ratio of oxygen to carbon (1.33:1 and 2.66:1)
○ for the same amount of carbon, one has exactly twice as much oxygen as the
other

● Dalton formulated his atomic theory in an attempt to explain how

and why elements would combine with one another in fixed ratios
and sometimes also in multiples of those ratios.
 THOMSON
• Experiments with cathode rays in 1897
• Cathode rays: streams of negatively charged
particles (electrons!)
→ led to the discovery of a fundamental building
block of matter – smaller than the smallest atom!

Cathode ray - Wikipedia

 THOMSON

• Negatively charged particles (electrons):

- Mass about 2,000 times smaller than a hydrogen atom
- Had same mass to charge ratio regardless of metals and gases used to
produce them
→ Electrons were found in all substances and were always the same.
 THOMSON’S ATOMIC THEORY
• Thomson also tried to show how the electrons were situated in the
atom.
• Since atoms were known to be electrically neutral, Thomson
proposed (1904) a model in which the atom was a positively
charged sphere studded with negatively charged electrons.
• It was called the “plum-pudding” model, since the electrons in the
atom resembled the raisins in a plum pudding.
 RUTHERFORD

• In 1911, experiments with alpha rays

→ A new atomic model: a small, heavy
nucleus with electrons in orbit around it.
→ This nuclear model of the atom became
the basis for the one that is still accepted
today.

Ernest Rutherford
(1871 - 1937)
 GOLD FOIL EXPERIMENT
• Over 98% of the
particles went straight
through

• About 2% of the
particles went through
but were deflected by
large angles

• About 0.01% of the

particles bounced off
the gold foil

http://micro.magnet.fsu.edu/electromag/java/rutherford/
GOLD FOIL EXPERIMENT (CONT.)
 RUTHERFORD’S ATOMIC THEORY
● The atom contains a tiny dense center called the nucleus
○ the volume is about 1/10 trillionth the volume of the atom
● The nucleus is essentially the entire mass of the atom
● The nucleus is positively charged
○ the amount of positive charge of the nucleus balances the
negative charge of the electrons
● The electrons move around in the empty space of the atom
surrounding the nucleus

AN ATOM IS
MOSTLY “EMPTY”
SPACE!
 PROTONS, ELECTRONS,
NEUTRONS
PARTICLE CHARGE MASS(u) MASS(kg) LOCATION
Proton (p) +1 1.0073 1.6725x10-27 Nucleus

Neutron (n) 0 1.0087 1.6750x10-27 Nucleus

Electron (e) -1 0.000549 9.1095x10-31 Moving around nucleus

If an atom were the size of a baseball stadium, Radius of Atom

the nucleus would be the size of a baseball at ------------------------ = 105
the center. Radius of Nucleus
 ATOMIC STRUCTURE
● Z = Atomic Number = number of protons in nucleus
○ For a neutral atom, number of protons = number of electrons

● A = Mass Number = sum of protons and neutrons in the nucleus

● Ions are atoms with a net charge due to gain or loss of electrons
○ Cations = positive = lost electrons
○ Anions = negative = gained electrons

● Given the atomic number, mass number, and the charge of an

atom, one can calculate the number of protons, electrons and
neutrons that make up that atom.
ISOTOPE NOTATION
Where:
A = mass number
Z = atomic number
X = element symbol

Carbon-12 has 6p, 6n, 6e

Oxygen-16 has 8p, 8n, 8e

Zinc-65 has 30p, 35n, 30e

 COUNTING SUBATOMIC
PARTICLES
O 8 8 8 8
Ruthenium 44 44 42 102

Iron Fe 26 30 3+
 ISOTOPES OF THE ELEMENTS
● Elements with the same atomic number but different
mass number
● Same number of protons, different number of neutrons
● Most elements have more than one isotope

Three isotopes of hydrogen

Three isotopes of neon
hydrogen

deuterium

tritium
AVERAGE ATOMIC MASS OR ATOMIC WEIGHT

● Weighted average of all (significant) naturally

occurring isotopes
● Takes into account mass and abundance of
each isotope

● Atomic Mass Unit (amu or u)

○ 1 u ~ 1.66 x 10-27 kg
 CALCULATING AVERAGE ATOMIC
MASS
Hydrogen-1 99.985%
Hydrogen-2 0.015% 0.99985 x 1 u + 0.00015 x 2 u = 1.00015 u

Relative Isotope
Oxygen-16 99.759% abundance mass
Oxygen-17 0.037%
Oxygen-18 0.204% Find the atomic mass of oxygen
and zinc based on the
Example for hydrogen
Zinc-64 48.89%
Zinc-66 27.81%
Zinc-67 4.11%
Zinc-68 18.57%
Zinc-70 0.62%
HOW IS RELATIVE ABUNDANCE DETERMINED?
Mass Spectroscopy
 NUCLEAR STABILITY AND DECAY:
RADIOACTIVITY
Stable isotopes do not break down over time
Radioactive isotopes are unstable

● Decay to produce new

nuclei
● Emit different radiation
depending on decay mode
● Emission of
particles/energy known as
radioactivity
 Nuclear Stability
Neutron:Proton (n/p) Ratio
● light elements (Z<20) n/p = 1:1
○ As Z increases, more neutrons are needed for stability
→ i.e. n/p increases
○ For heaviest stable isotopes, n/p => 1.5
● too high (too many neutrons)
○ beta decay
● too low (too many protons)
○ electron capture
○ positron emission
Mass
● too massive
○ alpha decay
Energy
● too much energy
○ gamma decay
 NUCLIDE CHART
● Plot of known nuclides (specific types of atoms or nuclei)
○ neutrons vs. protons

● Shows belt of stability

○ Region of stable isotopes
● This chart shows stable and unstable isotopes
Colorful Nuclide Chart
 Types of Radiation
● Alpha (α)

● Symbol: charge: +2 mass: 4 u

● Particles with two protons and two neutrons (Helium nucleus )
● Reduction of atomic number by two and mass number by four

● Beta (β)
● Symbol: charge: -1 mass: “0” (0.00055 u)
● High speed electron ejected from nucleus
● Negligible effect on mass number; increase atomic number by one

● Gamma (γ)
● Symbol: charge: 0 mass: 0
● High energy photons (bits of light)
● No effect on mass or charge
● Part of most nuclear reactions; release of pure energy
Types of Radiation (cont.)
 MODES OF RADIOACTIVE DECAY
● Alpha Decay
● Beta Decay
● Gamma Decay
● Electron Capture
● Positron Emission
Note: we will not consider the energetics or the emissions of other
particles like neutrinos which are required for conservation of energy
and momentum, so our discussion that follows is approximate.

We also will not consider half life at this time

 Alpha Decay
● Occurs for nuclei that are unstable due to large mass
number
● Nuclear reaction equations must conserve mass and
charge
● Every nucleus after Bi-83 is radioactive
 Beta Decay
● Occurs for nuclei whose n/p ratios are too high
● Involves decay of a neutron to give proton and
electron. Electron is ejected from nucleus.

● Atomic mass stays approx. the same, atomic number

increases by one
● Example
 Gamma Decay
● Energetic nucleus emits energetic gamma photon

● Usually associated with other nuclear transformations

● Nuclear charge and mass remains the same

→ No change in element
 Positron Emission
● occurs when there are too many
protons in the nucleus of an
atom
= n/p ratio too low

● Mass remains approx. the

same, atomic number
decreases by one
 Electron Capture
● occurs when (1) there are too
many protons in the nucleus of
an atom (= n/p ratio too low)
and
(2) have insufficient energy to
emit a positron

● Mass remains approx. the

same, atomic number
decreases by one
Decay Series for Uranium-238
 Nuclear Fission
● involves splitting a heavy nucleus into two or more
smaller, lighter ones.
● emits a significant amount of energy and neutrons
 Proton-proton chain
● One specific type of nuclear fusion (light atomic nuclei
combine to form heavier nuclei)
● Two or more atomic nuclei combine to form a larger
nuclei, nuclei/neutron by-products
● Also emits a significant amount of energy and neutrons
Summary – 3 types of
Nuclear Reaction
Summary of Atomic Models
The Periodic Table
 Periodic Table (cont.)
Vertical columns: groups or families
Horizontal rows: periods
Group A elements: main-group or representative elements
Group B elements: transition elements
two bottom rows: inner-transition (rare earth) elements

ALKALI METALS : I-A (Li, Na, K, Rb, Cs, Fr. Hydrogen is excluded; it is a nonmetal)
ALKALINE EARTHS (or alkaline earth metals) : II-A (Be, Mg, Ca, Sr, Ba, Ra)
HALOGENS : VII-A (F, Cl, Br, I, At)
NOBLE GASES (or inert gases) : VIII-A (He, Ne, Ar, Kr, Xe, Rn)
LANTHANIDES Elements 58 - 71
ACTINIDES Elements 90 - 103
 Metals vs. Non-Metals
Metals: tend to lose electrons when bonding with nonmetals
-form positive ions (cations)

Non-metals: tend to gain electrons when bonding with metals

-form negative ions (anions)
 ATOMS AND IONS
● Neutral atom: equal protons and electrons
● cation: electrons lost: positive charge
● anion: electrons gained: negative charge
Symbol Electrons Formula of Ion Name of Ion
Gained/Lost

Al 3 lost Al3+ aluminum ion

S 2 gained S2- sulfide ion
Ca 2 lost Ca2+ calcium ion
Br 1 gained Br- bromide ion
 COMPOUNDS
2 or more elements chemically combined
molecule: bonded group of atoms that act as unit
● molecular compounds
○ composed of molecules
○ two or more non-metals (e.g. CO2)
○ lower melting points
● ionic compounds
○ composed of positive and negative ions arranged in a
crystal lattice (3-d repeating array)
○ metal and non-metal (e.g. NaCl)
○ higher melting points
 CHEMICAL FORMULAS
● Show number/ratio of atoms of each element in smallest representative
unit of substance
● Molecular compounds
① Molecular formula
■ number of atoms of each element in one molecule
② Empirical Formula
■ shows simplest whole number ratio of elements in compound
③ Structural formula
■ Shows bonds between atoms with lines
④ Condensed structural formula
■ Arrangement shows important reactive groups

Acetic Acid (molecular)

Molecular formula Structural formula Condensed structural formula
C2H4O2 CH3COOH

Empirical Formula

CH2O
 CHEMICAL FORMULAS
● Show number/ratio of atoms of each element in smallest representative
unit of substance

● Ionic Compounds
○ Always shown as empirical formula
○ Formula unit
■ smallest whole number ratio of ions
■ e.g. the ratio of Na+ ions to Cl- ions in NaCl is 1:1

Calcium Fluoride (ionic)

Formula Unit: CaF2
 Writing Formulas Continued
Molecular
● Carbon is written first
● Hydrogen is second
● All others are written in alphabetical order
● If no carbon, alphabetical order for all
● For oxides, oxygen is last
● Examples: C6H12O6, CO2, CH4N2O, C3H7NO2S, P2O5
Ionic
● Positive ion first
● Negative ion second
● Net charge must be 0 (charges must balance!)
● Examples: NaCl, Al2O3, Ca(NO3)2, NH4Cl
Complete the Following Table in Your Notes
Structural Formula Molecular Formula Empirical Formula
 Writing Ionic Formulas
Write ionic compound formulas for the following:

SO42- Cl1- PO43-

NH41+

Al3+

Ca2+
 IONIC CHARGES OF THE ELEMENTS

● Common ionic charges of the representative

elements: based on group number
1A: +1, 2A: +2, 3A: +3
5A: -3, 6A: -2, 7A: -1
● Multivalent Atoms
○ Generally transition metals and others towards
bottom of p.table
○ have systematic and classical names
1+
2+ 3+ 3- 2- 1-
 NOMENCLATURE RULES
● Binary Ionic
○ representative elements (elements with one
common charge)
○ transition elements (multiple oxidation states)
● Binary Covalent
○ prefixes
● Ternary
○ polyatomic ions
● Acids
○ binary
○ ternary
Polyatomic Ions
Common Polyatomic Ions
● A polyatomic ion is a charged
group of covalently bonded
atoms.
● Common endings are -ate or -ite,
but there are exceptions.
● For more than 1 polyatomic ion, use
parentheses with the subscript on the
outside. There are 3 sulfate ions
Example: Al2(SO4)3 in this compound
 Polyatomic Ions
Sample Problem

Write chemical formulas for : Solution:

a. Calcium Hydroxide Ca2+1 OH 2– Ca(OH)2
b. Tin (IV) Sulfate Sn4+ SO4 2– Sn(SO4)2
2 4
Hint: Remember to divide subscripts by their largest common factor .
Write the correct names for:
a. (NH4)3 PO4 Ammonium Phosphate
2+ -
b. Cu(NO3)2 Cu (NO3)2 Copper(II) Nitrate
Hint: “Uncross” subscripts to get the charges of the ions.
 COMMON AND SYSTEMATIC NAMES

● Common names from the early discovery of

commonly used compounds
● systematic names follow rules that tell what the
compounds are made of
● examples:
○ K2CO3 "potash” potassium carbonate
○ CaO “lime” calcium oxide
○ KNO3 “saltpeter” potassium nitrate
○ NaOH “lye” sodium hydroxide
○ NaHCO3 baking soda sodium bicarbonate
Writing Binary Ionic Compounds
• Binary Compounds are composed of two
elements.
• Rules for writing binary ionic compounds:
1. Write the symbols for the ions, and their charges.
Note: The cation is always written first.
2. Cross over the charges (use the absolute value of each
ion’s charge as the subscript for the other ion.)
3. Simplify the numbers and remove the 1’s.
Example: aluminum oxide
Al23+ O32–
The correct formula for aluminum oxide Al
is O
2 3
Naming Binary Ionic Compounds

● The name of the cation is given first,

followed by the name of the anion.
● Monatomic cations are identified
simply by the element’s name.
● For monatomic anions, the ending
of the element’s name is dropped,
and the ending -ide is added.
Examples: cation anion
Al2O3 aluminum oxide
KF potassium fluoride
Binary Ionic Compounds
Sample Problem

Write chemical formulas for : Solution:

a. Magnesium Iodide Mg2+1
I –
2 MgI2
b. Calcium Oxide Ca22+ O 2–
2 CaO
Hint: Always divide subscripts by their largest common factor .
Write the correct names for:
a. Li2S Lithium Sulfide
b. ZnCl2 Zinc Chloride
Binary Ionic Compounds (m-nm) (elements with one common charge)

Roots for non-metals

H hyd
● Name of first element (positive ion) C carb
● name of second element + “ide” N nitr
ending (negative ion) P phosph
● examples: As arsen
● NaCl: sodium chloride O ox
● CaS: calcium sulfide S sulf
● Al2O3: aluminum oxide Se selen
F fluor
Cl chlor
Br Brom
I iod
 Multivalent Ionic Compounds
● Classical Name ION STOCK LATIN (classical)
● Same rules as before Cu+ Copper(I) Cuprous
● Use Latin name based Cu+2 Copper(II) Cupric
on charge Hg2+2 Mercury(I) Mercurous
● Stock Name Hg+2 Mercury(II) Mercuric
● Same rules as before Co+2 Cobalt(II) Cobaltous
● Use roman numeral to Co+3 Cobalt(III) Cobaltic
indicate charge on Mn+2 Manganese(II) Manganous
positive ion Mn+3 Manganese (III) Manganic
● Examples Fe+2 Iron(II) Ferrous
● FeCl2 iron(II) chloride or Fe+3 Iron(III) Ferric
ferrous chloride
Sn+2 Tin(II) Stannous
● CuSO4 copper(II) sulfate
Sn+4 Tin(IV) Stannic
or cupric sulfate
Pb+2 Lead(II) Plumbous
● Pb(NO3)2 lead(II) nitrate
or plumbous nitrate Pb+4 Lead(IV) Plumbic
The Stock System
● Most d-block elements (transition
metals) can form 2 or more ions
with different charges.
● To name ions of these elements,
scientists use the Stock system,
designed by Alfred Stock in 1919.
● The system uses Roman numerals to
indicate an ion’s charge.
Example: Fe2+ iron(II)
Fe3+ iron(III)
Stock System Naming
Sample Problem A

Write the formula and give the name for the

compound formed by the ions Cr3+ and F–.
Solution:
Cr 3+ F –
Write the ions side by side, cation first. 1 3
Cross over the charges to give subscripts. CrF3
Chromium forms more than one ion, so its
name must include the charge as a Roman
numeral.
Chromium (III) Fluoride
 Stock System Naming
Sample Problem B

Write chemical formulas for : Solution:

a. Tin (IV) Iodide Sn4+1 I – 4 SnI4
b. Iron (III) Oxide Fe23+ O 2–
3 Fe2O3

Write the correct names for:

3+ -
a. VF3 V F3 Vanadium (III) Fluoride
2+ 2-
b. CuO Cu O Copper (II) Oxide
Hint: “Uncross” subscripts to get the charges of the ions.
Be sure to verify the charge of the anion.
Classical name of transition metals
An older, less useful method for naming these
cations uses a root word with different suffixes
at the end of the word.

• The suffix -ous is used to name the cation

with the lower of the two ionic charges.

• The suffix -ic is used with the higher of

the two ionic charges.

• Using this system, Fe2+ is the ferrous

ion, and Fe3+ is the ferric ion.
A major disadvantage of using classical names
for ions is that they do not tell you the actual
charge of the ion.
 Nomenclature Flowchart
Compounds

Ionic Molecular Acids

Polyatomic Prefix Stock Binary Oxyacids

Binary System System Acids
Ions

Simple
(Main Group Elements)

Stock System
(d-Block Elements)
The Prefix System

● Molecular compounds are composed

of covalently-bonded molecules.
● The old prefix system is still used for
molecular compounds.
● Name the prefix, then the element.
Anions end in -ide.
● The prefix mono- usually isn’t used for the
first element.
Examples: P4O10 tetraphosphorus decoxide
CO carbon monoxide
The Prefix System
Sample Problem

Write chemical formulas for : Solution:

a. dinitrogen trioxide N2O3
b. carbon tetrabromide CBr4

Write the correct names for:

a. As2S3 diarsenic trisulfide
b. PCl5 phosphorus pentachloride
 TERNARY COMPOUNDS
● Generally involve polyatomic ion
● Name of first (positive) ion
● Name of second (negative) ion
● Examples:
○ NaHCO3 sodium bicarbonate
○ NH4Cl ammonium chloride
○ K2SO4 potassium sulfate
 Nomenclature Flowchart
Compounds

Ionic Molecular Acids

Polyatomic Prefix Stock Binary Oxyacids

Binary System System Acids
Ions

Simple
(Main Group Elements)

Stock System
(d-Block Elements)
Acids

• An acid is a certain type

of molecular compound.
All acids start with H
(e.g. HCl, H2SO4).
• Acids can be divided
into two categories:
1. Binary acids are acids that consist of H
and a non-metal. (e.g. HCl.)
2. Oxyacids are acids that contain H and a
polyatomic ion that includes O (e.g.
H2SO4.)
Binary Acids
• General rules for naming a binary acid:
1. Begin with the prefix hydro-.
2. Name the anion, but change the ending
to –ic.
3. Add acid to the name.
Examples:
HCl, hydrochloric acid.
HBr, hydrobromic acid.
H2S, hydrosulfuric acid.
Oxyacids
• General rules for naming an oxyacid :
1. Name the polyatomic ion.
2. Replace -ate with -ic or -ite with -ous
3. Add acid to the name.
Examples:
H2SO4, sulfuric acid.
H2SO3, sulfurous acid.
HNO3, nitric acid.
HNO2, nitrous acid.
Naming Acids
Sample Problem

Write the correct name for each of the following:

Type of Acid: Name:
a. HF binary acid hydrofluorine
ic acid
b. HNO2 oxyacid nitrite
ous acid
c. H2S binary acid hydrosulfuric acid
d. H2SO4 oxyacid sulfate
uric acid
e. H3PO4 oxyacid oric acid
phosphate
 SUMMARY OF NOMENCLATURE
● Be able to write name from formula or
formula from name
● Examples:
○ Sodium carbonate => Na2CO3
○ Dinitrogen tetroxide => N2O4
○ Hydrobromic acid => HBr
○ MgO => magnesium oxide
○ Fe3N2 => iron(II) nitride or ferrous nitride
 Nomenclature Flowchart
Compounds

Ionic Molecular Acids

Polyatomic Prefix Stock Binary Oxyacids

Binary System System Acids
Ions

Simple
(Main Group Elements)

Stock System
(d-Block Elements)