Chemistry: Modern Atomic Structure
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Atomic number: of an element is the number of protons of each atom of that element ** the
atomic number identifies an element**
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Because atoms are neutral, we know from the atomic number that there is the same number of
electrons.
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Atomic mass is the number of protons and neutrons in the nucleus (mass of the nucleus). To find
the number of neutrons in the nucleus of an atom you take the atomic mass and round it to the
nearest whole numbner then subtract the atomic number.
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All atoms of an element have the same number of protons. However, atoms of the same
element can have different number of neutrons. Isotope – atoms of the same element that
have different masses. Nuclide - general term for a specific isotope of an element.
Practice: 1. How many protons, electrons, and neutrons make up an atom of bromine-80?
2. Write the nuclear symbol for carbon-13. 3. Write the hyphen notation for the isotope with 15
electrons and 15 neutrons.
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Relative Atomic Masses: The standard by which scientists compare units of atomic mass if the
carbon-12 atom. It has been arbitrarily assigned a mass of exactly 12 atomic mass units, or 12
amu. 1 amu is exactly 1/12 the mass of a carbon-12 atom. Hydrogen-1 atom is 1.007825 amu.
An oxygen-16 atom has about 16/12 or 4/3 the mass of a carbon-12 atom. Oxygen-16 is
15.994915 amu. A magnesium-24 atom is found to be slightly les than twice that of a carbon-12
atom. Its atomic mass is 23.985042 amu.
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Masses of subatomic particles: electron = 0.0005486 amu, proton = 1.007276 amu, neutron =
1.008665 amu
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Average Atomic Masses of Elements: Most elements occur naturally as mixtures of isotopes.
The percentage of each isotope in the naturally occuring elements on Earth is nearly always the
same, no matter where the element is found. The percentage at which each of an elements
isotopes occurs in nature is taken into account when calculating the elements average atomic
mass. Average atomic mass is the weighted average at the atomic masses of the naturally
occuring isotopes of an element. * an elements atomic mass is sometimes rounded to two
decimal places before it is used in a calculation. *
Electron Configuration
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Electron configuration is the arrangement of electrons in an atom
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Aufbau principle – an electron occupies the lowest energy orbital that can receive it.
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Energy levels from lowest to highest: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f,
6d, 7p,
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S orbitals hold 2 e-, p orbitals hold 6 e-, d orbitals hold 10 e-, f orbitals hold 14eOrbital Notations
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Hund’s Rule – orbitals of equal energy are each occupied by one electron before any orbital is
occupied by a second electron, and all electrons in singly occupied orbitals must have the same
spin state.
The electron configuration of nitrogen is 1s22s22p3. How many electrons are present in a
nitrogen atom? What is the atomic number of nitrogen? Write the orbital notation for nitrogen.
The electron configuration of fluorine is 1s22s22p5. What is the atomic number of fluorine? How
many of its p orbitals are filled?
Write the electron configuration of sulfer.
Quantum Numbers & Atomic Orbitals
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In Bohr’s atomic model, electrons of increasing energy occupy orbitals farther and farther from
the nucleus.
Electrons in atomic orbitals also have quantized energies. There are also more characteristics of
orbitals.
The Four Quantum Numbers
1. Principle Quantum Number: symbolized by n, indicated the main energy level occupied by
the electron. Values of n are positive numbers 1-7. As n increases, the electron’s energy and its
average distance from the nucleus increase. More than one electron can have the same n value.
These electrons are said to be in the same electron shell.
2. Angular Momentum Quantum Number: Except at the first main energy level, orbitals of
different shapes, known as sublevels, exist for a given value of (n). The angular momentum
quantum number, symbolized by l (little L), indicates the shape of the orbital.
Shape
subshell
l
Ball
s
0
Dumbbell
p
1
4 pronged
d
2
8 lobed
f
3
3. Magnetic Quantum Number: symbolized by m, indicates the orientation of an orbital around
the nucleus.
Subshell
orientation
m
s
1
0
p
3
-1,0,1
d
5
-2,-1,0,1,2,
f
7
-3,-2,-1,0,1,2,3
4. Spin Quantum Number: symbolized by an s, has only two possible values (+1/2, -1/2) which
indicate the two fundamental spin states of an electron in an orbital. A single orbital can hold a
maximum of two electrons, which have opposite spin states.
Pauli Exclusion Principle: no two electrons in the same atom can have the same set of four
quantum numbers.
Diamagnetic – substance that contains no unpaired electrons
Noble Gas Notation
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The group 18 elements are called the noble gases. You can write shortened versions of electron
configuration using the noble gases.
Exceptions to the Electron Configurations
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There are a few elements that have slightly different electron configurations. These elements
have been proven to be more stable in an electron configuration that is slightly different from
the normal way we write our electron configurations.
Ions
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An ion is an element that has gained or lost electrons. Metals tend to lose electrons and have a
positive charge. Ions with a positive charge are called cations. Nonmetals tend to gain electrons
and have a negative charge. Ions with a negative charge are called anions. Oxidation numbers
are the numbers on the periodic table that tells us how many electrons an element is likely to
gain or lose. Electron configurations for ions will differ from the regular atom because the
electron number is different.
What is the electron configuration of Fe2+? How many protons, neutrons, and electrons are in
31
32 2
5
15P ? What elements ion could have the electron configuration of 1s 2s 2p ?
Electron Dot Notation
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Electron Dot Notation is an electron configuration notation in which only the valence electrons
of an atom of a particular element are shown, indicated by dots placed around the element’s
symbol.
The Octet Rule
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The Octet Rule: Unlike other atoms, the noble gas atoms exist independently in nature. This
stability results from the fact that, with the exception of helium and its two electrons in a
completely filled outer shell, the noble gas atom’s outer s and p orbitals are completely filled by
a total of eight electron, Other main group atoms can effectively fill their outermost s and p
orbitals with electrons by sharing electrons through covalent bonding. Chemical compounds
tend to form so that each atom, by gaining, losing, or sharing electrons, has eight electrons in its
highest occupied energy level.
Lewis Structures
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A Lewis structure is a representation of formulas in which atomic symbols represent nuclei and
inner-shell electrons, dot-pairs or dashes between two atomic symbols represent electron pairs
in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons.
It is common to write Lewis structures that show only the electrons that are shared, using
dashed to represent bonds. This is called a structural formula which indicates the kind, number,
arrangement, and bonds but not the unshared pairs of the atoms in a molecule. ** If carbon is
present in a compound, it is the central atom of the Lewis structure. **
Periodic Properties
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Atomic Radii: Ideally, the size of an atom is defined by the edge of its orbital. However, this
boundary is fuzzy and varies under certain conditions. One way to express an atom’s radius is to
measure the distance between the nuclei of two identical atoms that are chemically bonded
together, and then divided this distance by two. *****Atomic radii decreases as you go from left
to right across a period and increase as you go down a group******* There are some
exceptions!! Which of the following elements has the largest atomic radius: Li, O, C, or F? Which
has the smallest atomic radius?
Ionization Energy: (first ionization energies) An electron can be removed from an atom if
enough energy is supplied. Any process that results in the formation of an ion is referred to as
ionization energy (IE) – the energy required to remove one electron from a neutral atom of an
element. In general, ionization energies increase across a period and decrease down a group.
Group 1 metals have the lowest ionization energies in their respective periods. Therefore, they
lose electrons most easily. ** Low ionization energy means other atoms can easily take their
electrons. High ionization energies means that those atoms have enough energy to take
electrons from atoms.** This ease of electron loss is a major reason for the high reactivity of the
group 1 alkali metals. The low reactivity of noble gases is partly based on difficulty of electron
removal due to them having the highest ionization energy.
Electron Affinity: Neutral atoms can also acquire electrons. The energy change that occurs when
an electron is acquired by a neutral atom is called the atom’s electron affinity. Most atoms
release energy when they acquire an electron. On the other hand, some atoms must be “forced”
to gain an electron by the addition of energy. An ion produced in this way will be unstable and
will lose the added electron spontaneously. Among the elements of each period, the halogens
gain electron most readily. They have 7 valence electrons and only need one electron to become
stable. Trends for electron affinities within groups are not as regular as trends for ionization
energies. As a general rule, electrons add with greater difficulty going down a group.
Electronegativity: Valence electrons hold atoms together in chemical compounds. In many
compounds, the negative charge of the valence electrons is concentrated closer to one atom
than to another. This uneven concentration of charge has significant effect on the chemical
properties of a compound. It is therefore useful to have a measure of how strongly one atom
attracts the electrons of another atom with a compound. Electronegativity is a measure of the
ability of an atom in a chemical compound to attract electrons from another atom in the
compound. The most electronegative element is fluorine. Electronegativities tend to increase
across a period.
Absorption and Emissions
Electrons have the ability to move from one shell (energy level) (orbital) to another.
Emission – when an electron falls to a lower energy orbital. Higher number orbitals contain
more energy than those closer to the nucleus. As a result, when an electron moves down an
orbital it releases a photon that contains the energy difference between the two orbitals.
Absorption – moving an electron to a higher energy lever, this process requires energy to be
added.
Nuclear Reactions
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In nuclear chemistry, an atom is referred to as a nuclide and is identified by the number of
protons and neutrons in the nucleus. Most stable nuclei have a neutron to proton ratio of
approximately 1:1. As the atomic number increases, the ratio increases to 1.5:1. Beyond the
atomic number 83, bismuth, the repulsive force of the protons is so great that no stable nuclides
exist. Unstable nuclei undergo spontaneous changes that change their number of protons and
neutrons. In this process, they give off large amounts of energy and increase their stability.
These changes are a type of nuclear reaction. In equations representing nuclear reactions, the
total of the atomic numbers, and the total of the mass numbers must be equal on both sides of
the equation. Transmutation – change in the identity of a nucleus as a result of a change in the
number of its protons.
Radioactive Decay
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Radioactive decay is the spontaneous disintegration of a nucleus into a slightly lighter nucleus,
accompanied by the emission of particles or electromagnetic radiation, or both
Particles Chart
Alpha Emission (α) is two protons and two neutrons bound together and are emitted from the
nucleus during some kinds of radioactive decay. Alpha emission is restricted almost entirely to
very heavy nuclei.
Beta Emissions (β): To decrease the number of neutrons, a neutron can be converted into a
proton and an electron. A beta particle is an electron emitted from the nucleus during some
kinds of radioactive decay.
Gamma Emission ( ): Gamma rays are high-energy electromagnetic waves emitted from a
nucleus as it changes from an excited state to a ground energy state. According to the nuclear
shell model, gamma rays are produced when nuclear particles undergo transitions in nuclearenergy levels. Gamma emission usually occurs immediately following other types of decay.
Positron Emission: To decrease the number of protons, a proton can be converted into a
neutron by emitting a positron. A positron is a particle that has the same mass as an electron,
but has a positive charge, and is emitted from the nucleus during some kinds of radioactive
decay.
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Half-Life
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No two radioactive isotopes decay at the same rate. Half-life, t1/2 , is the time required for half
the atoms of a radioactive nuclide(nucleus) to decay. More stable nuclides decay slowly and
have longer half-lives. Less stable nuclide decay very quickly and have shorter half-lives.
Sometimes just a fraction of a second.
The half-life of polonium-210 is 138.4 days. How many milligrams of polonium-210 remain after
415.2 days if you start with 2.0mg of the isotope?
Assuming a half-life of 1599 years, how many years will be needed for the decay of 15/16 of a
given amount of radium-226?
The half-life of radon-222 is 3.824 days. After what time will one-fourth of a given amount of
radon remain?
The half-life of cobalt-60 is 5.27 years. How many milligrams of cobalt-60 remain after 52.7
years if you start with 10.0 mg?