Grade 11 Physical Science Atomic Theory

Atomic Number
The atomic number of an element is determined by the number of protons in it, and it is
used to differentiate one element from another. The mass number of an element is
determined by the number of protons and neutrons combined.
The modern periodic table is arranged in such a way that all the elements have an
increasing atomic number, and subsequently, increasing mass number. But do you
know what mass number, or even what atomic number is?
Well, as you know, an atom consists of electrons, protons and neutrons. The number of
electrons in the outermost shell gives us the valency of the atom. Similarly, the number of
protons and neutrons are associated with the atomic number and mass number of the
atom.

What is Atomic Number?
● The total number of protons in the nucleus of an atom gives us the
atomic number of that atom.
● It is represented with the letter ‘Z.’
● All the atoms of a particular element have the same number of
protons, and hence the same atomic number.
● Atoms of different elements have different atomic numbers.
● For example, all carbon atoms have the atomic number of 6,
whereas all atoms of Oxygen have 8 protons in their nucleus.

What is Mass Number?
● The number of protons and neutrons combine to give us the mass number of
an atom.
● It is represented using the letter ‘A.’
● As both protons and neutrons are present in the nucleus of an atom, they are
together called nucleons.
● For example, an atom of carbon has 6 protons and 6 neutrons. Thus, its
mass number is 12.
● While the number of protons remains the same in all atoms of an element,
the number of neutrons can vary. Thus, atoms of the same element can have
different mass numbers, and these are called isotopes.
● The weight of an electron is almost negligible. Thus, the atomic mass of an
atom is almost the same as its mass number.

Isotopes are atoms with the same atomic number but distinct
neutron numbers, and hence distinct mass numbers.
● The average isotopic mass of an isotopic mixture for an
element in a defined environment on Earth determines the
element’s standard atomic weight.
● A little more than three-quarters of naturally occurring
elements exist as a mixture of isotopes, and the average
isotopic mass of an isotopic mixture for an element in a
defined environment on Earth determines the element’s
standard atomic weight.

Examples of Atomic Number
The atomic number of an atom is equal to the number of protons in the nucleus of an atom or the number of
electrons in an electrically neutral atom.
Atomic number = Number of protons
For example, in a sodium atom, there are 11 electrons and 11 protons. Thus the atomic number of Na atom =
number of electrons = number of protons = 11.

Atomic Number Orbital
Energy Levels
When an electron is at a specific energy
level, it is more likely to be found in
certain portions of that level than others.
Orbitals are the name for these sections.
Sublevels are made up of orbitals with
the same energy. A maximum of two
electrons can be found in each orbital.
The most common way of showing the
arrangement of electrons in an atom is
to draw diagrams like those shown in the
diagram.

Atoms Are Neutral
The northern lights aren’t caused by atoms, because atoms are not charged particles. An atom always has
the same number of electrons as protons. Electrons have an electric charge of -1 and protons have an
electric charge of +1. Therefore, the charges of an atom’s electrons and protons “cancel out.” This explains
why atoms are neutral in electric charge.
Q: What would happen to an atom’s charge if it were to gain extra electrons?
A: If an atom were to gain extra electrons, it would have more electrons than protons. This would give it a
negative charge, so it would no longer be neutral.

Atoms to Ions
Atoms cannot only gain extra electrons. They
can also lose electrons. In either case, they
become ions. Ions are atoms that have a
positive or negative charge because they
have unequal numbers of protons and
electrons. If atoms lose electrons, they
become positive ions, or cations. If atoms
gain electrons, they become negative ions,
or anions. Consider the example of fluorine
(see Figure below). A fluorine atom has nine
protons and nine electrons, so it is
electrically neutral. If a fluorine atom gains an
electron, it becomes a fluoride ion with an
electric charge of -1.

Cations Anions
Name of Ion Chemical Symbol Name of Ion Chemical Symbol
Calcium ion Ca2+
Chloride Cl-
Hydrogen ion H+
Fluoride F-
Iron(II) ion Fe2+
Bromide Br-
Iron(III) ion Fe3+
Oxide O2-
just given the element name followed
by the word ion. For example, when
a sodium atom loses an electron, it
becomes a positive sodium ion. The
charge of an ion is indicated by a
plus (+) or minus sign (-), which is
written to the right of and just above
the ion’s chemical symbol. For
example, the fluoride ion is
represented by the symbol F-
, and
the sodium ion is represented by the
symbol Na+
. If the charge is greater
than one, a number is used to
indicate it. For example, iron (Fe)
may lose two electrons to form an
ion with a charge of plus two. This
ion would be represented by the
symbol Fe2+
. This and some other
common ions are listed with their
symbols in the Table below.

How Ions Form
The process in which an atom
becomes an ion is called ionization.
It may occur when atoms are
exposed to high levels of radiation.
The radiation may give their outer
electrons enough energy to escape
from the attraction of the positive
nucleus. However, most ions form
when atoms transfer electrons to or
from other atoms or molecules. For
example, sodium atoms may transfer
electrons to chlorine atoms. This
forms positive sodium ions (Na+
) and
negative chloride ions (Cl-
).

Why do you think atoms lose
electrons to, or gain electrons from,
other atoms?
A: Atoms form ions by losing or
gaining electrons because it makes
them more stable and this state
takes less energy to maintain. The
most stable state for an atom is to
have its outermost energy level filled
with the maximum possible number
of electrons. In the case of metals
such as lithium, with just one
electron in the outermost energy
level, a more stable state can be
achieved by losing that one outer
electron. In the case of nonmetals
such as fluorine, which has seven
electrons in the outermost energy
level, a more stable state can be
achieved by gaining one electron

Properties of Ions
Ions are highly reactive, especially
as gases. They usually react with
ions of opposite charge to form
neutral compounds. For example,
positive sodium ions and negative
chloride ions react to form the neutral
compound sodium chloride,
commonly known as table salt. This
occurs because oppositely charged
ions attract each other. Ions with the
same charge, on the other hand,
repel each other. Ions are also
deflected by a magnetic field, as you
saw in the opening image of the
northern lights.

What are Electron Configurations?
The electron configuration of an element describes how electrons are distributed in its atomic orbitals. Electron
configurations of atoms follow a standard notation in which all electron-containing atomic subshells (with the number of
electrons they hold written in superscript) are placed in a sequence. For example, the electron configuration of sodium is
1s22s22p63s1.

However, the standard notation often yields lengthy electron configurations (especially for elements
having a relatively large atomic number). In such cases, an abbreviated or condensed notation may be
used instead of the standard notation. In the abbreviated notation, the sequence of completely filled
subshells that correspond to the electronic configuration of a noble gas is replaced with the symbol of
that noble gas in square brackets. Therefore, the abbreviated electron configuration of sodium is
[Ne]3s1 (the electron configuration of neon is 1s22s22p6, which can be abbreviated to [He]2s22p6).

Electron Configurations are useful for:
● Determining the valency of an element.
● Predicting the properties of a group of elements (elements with similar electron configurations tend to exhibit
similar properties).
● Interpreting atomic spectra.
This notation for the distribution of electrons in the atomic orbitals of atoms came into practice shortly
after the Bohr model of the atom was presented by Ernest Rutherford and Niels Bohr in the year 1913.

Shell and ‘n’ value Maximum electrons present in the shell
K shell, n=1 2*12 = 2
L shell, n=2 2*22 = 8
M shell, n=3 2*32 = 18
N shell, n=4 2*42 = 32
Writing Electron
Configurations
Shells
The maximum number of electrons that can
be accommodated in a shell is based on the
principal quantum number (n). It is
represented by the formula 2n2, where ‘n’ is
the shell number. The shells, values of n, and
the total number of electrons that can be
accommodated are tabulated below.

Subshells
● The subshells into which electrons are distributed are based on the azimuthal quantum number (denoted by ‘l’).
● This quantum number is dependent on the value of the principal quantum number, n. Therefore, when n has a value of 4, four different
subshells are possible.
● When n=4. The subshells correspond to l=0, l=1, l=2, and l=3 and are named the s, p, d, and f subshells, respectively.
● The maximum number of electrons that can be accommodated by a subshell is given by the formula 2*(2l + 1).
● Therefore, the s, p, d, and f subshells can accommodate a maximum of 2, 6, 10, and 14 electrons, respectively.

Aufbau Principle
● This principle is named after the German word ‘Aufbeen’ which means ‘build up’.
● The Aufbau principle dictates that electrons will occupy the orbitals having lower energies before occupying higher energy
orbitals.
● The energy of an orbital is calculated by the sum of the principal and the azimuthal quantum numbers.
● According to this principle, electrons are filled in the following order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p,
7s, 5f, 6d, 7p…

Pauli Exclusion Principle
● The Pauli exclusion principle states that a maximum of two electrons, each having opposite spins, can fit in an orbital.
● This principle can also be stated as “no two electrons in the same atom have the same values for all four quantum
numbers”.
● Therefore, if the principal, azimuthal, and magnetic numbers are the same for two electrons, they must have opposite
spins.
Hund’s Rule
● This rule describes the order in which electrons are filled in all the orbitals belonging to a subshell.
● It states that every orbital in a given subshell is singly occupied by electrons before a second electron is filled in an orbital.
● In order to maximize the total spin, the electrons in the orbitals that only contain one electron all have the same spin (or
the same values of the spin quantum number).

Electron Configuration of
Hydrogen
The atomic number of hydrogen is 1.
Therefore, a hydrogen atom contains 1
electron, which will be placed in the s
subshell of the first shell/orbit. The
electron configuration of hydrogen is 1s1,
as illustrated below.

Electron Configuration of
Oxygen
The atomic number of oxygen is 8,
implying that an oxygen atom holds 8
electrons. Its electrons are filled in the
following order:
K shell – 2 electrons
L shell – 6 electrons
Therefore, the electron configuration of
oxygen is 1s2 2s2 2p4, as shown in the
illustration provided below.

Chlorine Electronic Configuration
Chlorine has an atomic number of 17. Therefore, its 17
electrons are distributed in the following manner:
K shell – 2 electrons
L shell – 8 electrons
M shell – 7 electrons
The electron configuration of chlorine is illustrated
below. It can be written as 1s22s22p63s23p5 or as
[Ne]3s23p5

What is Electron Dot Structure?
Electron dot structures or Lewis dot formula can be drawn if the molecular formula of the compound is known.
It defines the nature of bond and position of atoms of the molecule which are connected in the molecule.
The representation of molecules in Lewis electron dot structure or just a Lewis structure is in honour of the American
chemist Gilbert Newton Lewis.
Lewis dot structures also called electron dot structures are diagrams that describe the chemical bonding between
atoms in a molecule. They also display the total number of lone pairs present in each of the atoms that constitute the
molecule. Lewis dot structures are commonly referred to as electron dot structures or Lewis structures. Lewis
defined a base as an electron pair donor and an acid as an electron pair acceptor.

Lewis Dot Structure
Lewis dot structures reflect the electronic structures of the elements, including how the electrons are paired.
Lewis structures are a useful way to summarize certain information about bonding and may be thought of as
“electron bookkeeping”. In Lewis dot structures each dot represents an electron. A pair of dots between chemical
symbols for atoms represents a bond.

Imagine you’re organizing a concert, and
each ticket has a specific code to tell people
where they should sit. In the world of atoms,
"quantum numbers" are like these ticket
codes—they tell us where electrons "sit" or
"exist" around the nucleus of an atom.
There are four quantum numbers, and
each one gives specific information about the
"seat" or "state" of an electron:

1. Principal Quantum Number (n)
● What it tells us: How far the electron
is from the nucleus.
● Think of it as: The row number in a
concert hall.
● Values: Whole numbers like 1, 2, 3,
etc.
● Example: If n=1n = 1n=1, the
electron is in the closest "seat" to the
nucleus. If n=3n = 3n=3, it’s farther
away.

2. Angular Momentum Quantum
Number (l)
● What it tells us: The shape of the
space (or "seat section") where the
electron is likely to be found.
● Think of it as: The type of seating
section—like orchestra, balcony, or
mezzanine.
● Values: Starts at 0 and goes up to
n−1n-1n−1.
○ l=0l = 0l=0: Sphere-shaped
(called "s").
○ l=1l = 1l=1: Dumbbell-shaped
(called "p").
○ l=2l = 2l=2: Clover-shaped
(called "d").
○ l=3l = 3l=3: More complex
shapes (called "f").

3. Magnetic Quantum Number (mĻ)
● What it tells us: The orientation of
the "seat section" in 3D space.
● Think of it as: Which exact seat
within the section.
● Values: From −l-l−l to +l+l+l, including
0.
○ For l=1l = 1l=1: mlm_lmlcould
be −1,0,+1-1, 0, +1−1,0,+1,
representing different
dumbbell orientations.

4. Spin Quantum Number (ms_ss
)
● What it tells us: The "spin" of the electron, which is a tiny magnetic property.
● Think of it as: Whether the person in the seat is clapping up or down.
● Values: +12+frac{1}{2}+21(spin up) or −12-frac{1}{2}−21(spin down).

Grade 11 Physical Science Atomic Theory

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