VIETNAM NATIONAL UNIVERSITY, HO CHI MINH CITY
INTERNATIONAL UNIVERSITY
WEEK 3
ATOMS, MOLECULES AND IONS
Part 02
S2_2023-2024
References:
(1) Chemistry by John E. McMurry, Robert C. Fay, Jill K. Robinson, Chapter 5, 7th
Eds., published by Pearson Education@2015.
Instructor: Dr. Ngo Thi Thuan
12/03/2024
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Learning outcomes
STUDENTS SHOULD BE
ABLE TO…
G1 Write electron configuration of atoms
G2 Define an atom’s group and period in the periodic table
from electron configurations
G3 Clarify ions and their charges from electron
configurations
G4 Distinguish between molecules and ionic compounds in
terms of chemical bonding
G5 Give chemical formula and names of molecules and ionic
compounds, acids and simple organics
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1. Introduction
Group
(column)
Period (row)
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• Main group: Group number = number of
valence electron = electrons at the
outermost shell
• Period number = number of energy
levels= number of orbits
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• Electron number at the outermost shell
= valence shell containing the valence
4 valence electrons
electrons = main group number.
Example:
& 2 core electrons
6 protons • Period number = number of
6 neutron
energy levels (shells).
Group 4A
Period 2
Electron
Proton
Neutron
Carbon atom
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2. Quantum numbers
A set of numbers used to describe the position and
energy of the electrons in an atom
Principal quantum number:
energy level (n)
Angular momentum:
Sublevel energy(l)
Magnetic number:
Orbital numbers in sublevel (ml)
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Spin number:
Electron spins (ms)
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❖Energy level (n) (principal quantum number): a
regions where electron are referred
• Electrons in the outermost
circles have higher energy.
• The
energy
levels
are
numbered 1, 2, 3…and
represented as letter “n”.
❖ Sublevel of energy (l): Within each energy level are
sublevel, which depict general shape and well-defined
energy of the electron cloud (orbital).
Electron cloud (orbital): a region where there is high
probability of finding the electrons.
Example: electron distribution in the ground state (n=1) of the
hydrogen atom (Bohr model)
Schrodinger
Equation
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❖ Sublevel of energy (l)
Name of orbitals
Sublevel (l)
s
Spherical
0
p
Dumbbell
1
d
Cloverleafs
2
f
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9
➢ Sublevel (l) can have any integral value, starting with 0
and a maximum of (n-1):
ll = 0, 1, 2, …,(n-1)
n
1
2
Sublevel (l) 0 0
1 0
1s 2s
2p 3s
3
4
1
2 0 1
2
3p 3d 4s 4p 4d
3
4f
➢ Sublevel increases energy in order of:
ns <np < nd < nf
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❖ Orbital number in sublevel (ml): Within each sublevel,
there are orbitals which describe direction in space of the
electron cloud surrounding the nucleus.
s
l = 0, ml = 0 →1 orbital
p
l = 1, ml = -1, 0, +1
→ 3 orbitals
d
f
l = 2,
ml = -2, -1, 0, +1, +2
→ 5 orbitals
l = 3,
ml = -3,-2, -1, 0, +1, +2, +3
→ 7 orbitals
❖ ml can have any integral value, including 0, between l
and – l
ml = l, …,+1, 0, -1, …, -l
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❖ Electron spin (ms): describe the spin orientation of
electron in 01 orbital
➢
A maximum of 02 electrons occupy 01 orbital and spin in opposite direction
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❖ Electron spin (ms): describe the spin orientation of
electron in 01 orbital
Pauli Exclusion Principle: “No two electrons in an atom can have
same set of four quantum number” → two electrons in an atom can
have same n, l and m value but differ in spin number
Example: He (Z =2); n =1, l = 0 , ml = 0 but ms = -1/2, ms =+1/2
➢
A maximum of 02 electrons occupy 01 orbital and spin in opposite direction
➢ Useful to determine the max no. of e- in:
+ Main energy level (n) = 2n2
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➢ Maximum of electron number in sublevels
s
2 electrons
p
6 electrons
d
10 electrons
f
14 electrons
2(2l + 1)
Summary of quantum numbers
Names of quantum Symbol
numbers
Values
Properties
Principal quantum
number (level
energy)
n
Positive
Orbital energy
integers (1, 2,
3, …)
Angular momentum
(sublevel energy)
l
Magnetic number
(Orbital number)
ml
Integer from Orbital shape
0 to n-1
l=0: s orbital
l=1: p orbital
l=2: d orbital
l=3: f orbital
Integers from Orbital direction
–l to 0 to + l
Electron spins
ms
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+1/2 or -1/2
Direction of 2 ein 01 orbitals 15
Example: Determine electron distribution of Li atom
(Z = 3)
-
n
l
ml
ms
e distribution
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1s
1
0
0

1s2
2s
2
0
0

2s1
2p
2
1
-1, 0, +1
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2. Electron configuration: electron distribution among
available subshell
Aufbau principle
Q&A: How do we arrange
electrons into these orbitals?
Hund’s rule
Aws: Increasing energy
➢ Putting electrons into separate orbitals of the subshell with the
same spin before pairing electrons.
Example
Write electron configuration
of Se and define its position in the periodic table?
Period 4 and 6A
IE 1 (2 mins)
Write electron configuration
of Al and define its position in the periodic table?
Aufbau Principle
1s2 2s2 2p6 3s2 3p1
Period 3, 3A
Electron configuration_main and d- groups
Main-group elements (A)
(valence electron: nsa npb)
s-block elements
p-block elements
Transition elements: the other elements
Periodic table _Electron configuration at ground state)
Atomic
number
Partial Orbital
Diagram
Element
3s
11
Na
12
Mg
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
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Full Electron
Configuration
Condensed Electron
configuration
3p
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Electron configuration _ Exception to the Aufbau Principle
Half-full or full shells of subshells are more stable than partially
filled ones.
Example:
Copper: 29 electrons, how to fill them into orbitals?
Aufbau Principle 1s2 2s2 2p6 3s2 3p6 4s2 3d9
Exception 1s2 2s2 2p6 3s2 3p6 4s1 3d10
This electron configuration is more stable
because of half and full valence subshells
Full valence shell = low chemical reactivity
[Ar] 4s1 3d10
Period 4, 1B (d-block)
IE
Write electron configuration of Chromium
and define its position in the periodic table?
Aufbau Principle 1s2 2s2 2p6 3s2 3p6 4s2 3d4
Exception 1s2 2s2 2p6 3s2 3p6 4s1 3d5
This electron configuration is more stable
because of half-full valence shell
[Ar] 4s1 3d5 Period 4, 6B (d-block)
3. Ions
Atom loses or
gains electrons,
it becomes an
ion.
Predicting ion charge of ions in main group
Q&A : How would an atom want
to lose or gain electrons?
Feel like “noble gas”
2s22p2
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2s22p3
2s22p4
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Predicting charge on d-block ions
Elements in the d-block (called the transition metals) can exist as
ions with various charges due to various numbers of electrons in
d-orbitals and depend on what they’re bonded to.
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Predicting charge on d-block ions
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https://chem.libretexts.org/
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Charge of polyatomic ions
• A group of atoms will gain or lose electrons. These are
polyatomic ions.
Ex: A polyatomic cation: NH4+
A polyatomic anion: SO42• Overall charge of polyatomic ions can be determined
via charge of each element
Example:
SO42- (S (+6) and O (-24=-8) = -2)
NO3- (N (+5) and O (-23 =-6) = -1)
NH4+ (N(-5) and H (+14=+4) = +1)
Electron numbers in ions
• Number of electrons in ion = number of protons
(atomic number) – its charge
Example: how many electrons in 131I- are there?
4. Compound
- Atoms form bond to fill the outer shell with
electrons →compound
Sharing
electrons
among atoms (covalent
bond) →molecular (ex.
H2O, Cl2….). Usually
nonmetals
Transferring
electrons
from this atoms to those
of another (ionic bond)→
ionic compounds (ex.
NaCl). Usually metal and
nonmetal
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COMPOUND
Covalent bond
A molecule
(covalent compounds)
Ionic bond
A formula unit
(ionic compounds)
The lowest whole number ratio of
ions in an ionic compound
Some diatomic Molecules
• These seven elements occur naturally as molecules containing
two atoms:
– Hydrogen (H2)
– Nitrogen (N2)
– Oxygen (O2)
– Fluorine (F2)
– Chlorine (Cl2)
– Bromine (Br2)
– Iodine (I2)
Some ionic compounds
Note: Ionic compounds are made up of a metal and a non-metal
(except ammonium chloride which is made up only non-metals
5. Chemical Formulas
- To represent compound
Symbols of elements
Relative number of
atoms (subscript)
Writing compound formulas
• Charge of compounds are electrically neutral
(principle of charge balance)
+ The charge on the cation becomes the subscript on
the anion.
+ The charge on the anion becomes the subscript on
the cation.
+ If these subscripts are not in the lowest wholenumber ratio, divide them by the greatest common
factor.
6. Chemical nomenclature
• The system of naming compounds is called
chemical nomenclature.
• We will learn how to name:
1) Ionic compounds vs. molecular
2) Acids
3) Simple Organic Compounds
–Alkanes
–Alcohols
6.1. Ionic compounds (Type I_ main group metals)
•The cation (metal) is always named first without
Roman numeral of cation charge
•The anion (nonmetal) is written after the cation,
modified to end in –ide
Compound
Ions
present
Name
NaCl
Na+
Cl-
Sodium chloride
KI
K+
I-
Potassium iodide
CaS
Ca2+ S2-
Calcium sulfide
Li3N
Li+
N3+
Lithium nitride
CsBr
Cs+
Br-
Cesium bromide
MgO
Mg2+ O2-
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Magnesium oxide
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Inorganic Nomenclature_main group metals
+1 Charge
+2 charge
-1 Charge
Group 1A
elements
Group 2A
elements
Hydrogen: H+
Beryllium: Be2+
Lithium: Li+
Magnesium: Mg2+ Fluoride: F-
Sodium: Na+
Calcium: Ca2+
Chloride: Cl-
Potassium: K+
Strontium: Sr2+
Bromide: Br-
Rubidium: Rb+ Barium: Ba2+
Cesium: Cs+
Group 7A
elements
Iodide: I-
IE
•Given the following systematic names, write the
formula for each compound:
a. Potassium iodide
b. Calcium oxide
c. Gallium bromide
Solution:
Name
Formula
Comments
Potassium iodide
KI
Contains K+ and I-
Calcium oxide
CaO
Contains Ca2+ and O2-
Gallium bromide
GaBr3
Contains 3Br- to balance
charge if Ga3+
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6.2. Ionic compounds (Type II_transition metals)
Ion
Systematic Name
Fe3+
Iron (III)
Fe2+
Iron(II)
Cu+
Copper(I)
Cu2+
Copper(II)
Co3+
Cobalt(III)
Co2+
Cobalt(II)
Sn4+
Tin(IV)
Sn2+
Tin(II)
Pb2+
Lead(II)
Pb4+
Lead(IV)
Hg2+
Mercury(II)
2+
Mercury (I)
Hg2
Ag+
Silver
Zn2+
Zinc
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Cd2+
Cadmiun
•The cation (metal) is always named
first with Roman number of cation
charge.
•The anion (nonmetal) is written after
the cation, modified to end in –ide
Ex.: CuCl → Copper (I) chloride
Mercury (I) ions always occur
bound together to form Hg22+
ions (ex. Hg2Cl2 →Mercury(I)
chloride)
Although these are transition
metals, they form only one type
of ion and a Roman numeral is
not used
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IE4 (3 mins)
•Give the systematic names for each compounds
a. HgO
b. AgCl
c. Fe2O3
Solution:
Formula
Name
Comments
HgO
Mercury(II) oxide
The anion is O2- , the cation must
be Hg2+
AgCl
Silver chloride
The anion is Cl- , silver has only
one type of Ag+ , Roman numeral
is not used
Fe2O3
Iron (III) oxide
The three O2- carry a total charge
of 6-, so two Fe3+ ions are needed
to give a 6+ charge
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Ionic compounds (Type II_transition metals)_another
system
• Cation (metal) is always named first (-ous ending for the lower
oxidation state; -ic ending for the higher oxidation state)
• Anion (nonmetal) is written after the cation, modified to end
in –ide
Example:
FeCl3 : ferric chloride
FeCl2 : ferrous chloride
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Does the ionic
compound contain
Type I or Type II
cations?
Type I cations
Name the
cation using
the element
name
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Type II cation
Using the principal of
charge balance →
determine the cation
charge.
Include in the cation name,
a Roman numeral
indicating the charge of
cation, except for Ag, Zn
and Cd
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6.4. Ionic compounds with polyatomic ions
Patterns of Oxyanion ions
Anions contain an atom of a given element and different
numbers of oxygen atoms
Maximum of
three
O atoms in
period 2
Period▪2
Charges
increase from
right to left
▪
Period 3
Maximum of
four
O atoms in
period 3
Carbon
Oxygen
Names of Oxyanion ions
• When there are more two oxyanions involving the same
element
– the one with fewer oxygens ends in -ite.
– the one with more oxygens ends in -ate.
– hypo-(less than) and per-(more than) are used as prefixes to
name the members with the fewest and the most oxygen
atoms, respectively
IE5 (3 mins)
•Give the systematic names for each compounds
a. Na2SO4
b. Fe(NO3)3
c. NaClO4
Solution:
Formula
Na2SO4
Fe(NO3)3
NaClO4
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Name
Sodium sulfate
Iron (III) nitrate or ferrous
chloride
Sodium perchlorate
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6.5. Naming molecular compounds
The first element is named first, using the full
element name
The second element is named, The ending
on the second element is changed to -ide
Using prefix to denote the numbers of atoms
present
Never use prefix mono- for naming the first
element. Two successive vowels are often
elided into one.
N 2O
Systematic Name
Dinitrogen monoxide
NO
Nitrogen monoxide
Common Name
Nitrous oxide
Nitric oxide
IE6 (3 mins)
•Give the systematic names for each compounds
a. CO2
b. CO
c. CCl4
d. N2O5
Solution:
Formula Name
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CO2
carbon dioxide
CO
carbon monoxide
CCl4
carbon tetrachloride
N2O5
dinitrogen pentoxide (pentaoxide)
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6.6. Acid Nomenclature
• If the anion in the acid ends in ide, change the ending to -ic
acid and add the prefix hydro-.
– HCl: hydrochloric acid
– HBr: hydrobromic acid
– HI: hydroiodic acid
• If the anion ends in -ite, change the ending to -ous acid.
– HClO: hypochlorous acid
– HClO2: chlorous acid
• If the anion ends in -ate, change the ending to -ic acid.
– HClO3: chloric acid
– HClO4: perchloric acid
6.7. Nomenclature of Organic
Compounds: Alkanes
• Organic chemistry is the study of carbon.
• Organic chemistry has its own system of nomenclature.
• The simplest hydrocarbons (compounds containing only carbon
and hydrogen) are alkanes.
• The first part of the names just listed correspond to the number
of carbons (meth- = 1, eth- = 2, prop- = 3, etc.).
• It is followed by -ane.
Nomenclature of Organic Compounds:
Alcohols (1 of 2)
• When a hydrogen in an alkane is replaced with
something else (a functional group, like −OH in the
compounds above), the name is derived from the
name of the alkane.
• The ending denotes the type of compound.
– An alcohol ends in -ol.
Name of common Cations
Common Cations
aThe
ions we use most often in this course are in boldface. Learn them first.
Common Anions
a The
ions we use most often are in boldface. Learn them first.