Download Identify the relationships among the components of the atom

Identify the relationships among the components of the atom.
Learning Focus:
Identify protons, neutrons, and electrons as components of atoms
Consider the forces which hold the atom together.
Recognize the terminology used to describe atoms and their isotopes: atomic number; nucleon
(mass) number, atomic mass, atomic mass unit average atomic mass.
Distinguish between isotopes of an element.
Recognize that there is a difference between mass and weight.
Atom – the smallest particle of an element that retains the properties of the element.
electron cloud
nucleus
proton
 If an atom had a
diameter of two
football fields, the
nucleus would be
the size of a nickel.
neutron
Atoms are spherically shaped, with a tiny, dense nucleus of positive charge
surrounded by one or more negatively charged electrons. Most of an atom
consists of fast-moving electrons traveling through the empty space surrounding
the nucleus. The electrons are held within the atom by their attraction to the
positively charged nucleus. The nucleus, which is composed of neutral neutrons
(hydrogen’s single-proton is an exception) and positively charged protons,
contains all of the atom’s positive charge and more than 99.97% of its mass.
Since an atom is electrically neutral, the number of protons in the nucleus equals
the number of electrons surrounding the nucleus.
 If a nucleus were
the size of the dot in
the exclamation
point at the end of
this sentence, its
mass would be
approximately as
much as that of 70
automobiles!
Properties of Subatomic Particles
Particle
Symbol
-
Electron
e
Proton
p+
Neutron
no
Location
In the space
surrounding
the nucleus
In the
nucleus
In the
nucleus
Relative
electrical
charge
Actual mass
(g)
Mass
(amu)
Relative
mass
1-
1/1840
9.11x10-28
0.000549
1+
1
1.673x10-24
1.007276
0
1
1.675x10-24
1.008665
A periodic table has over 110 different elements. This means that there are more than
110 different kinds of atoms. Atoms of each element contain a unique positive charge in
their nuclei. Therefore, the number of protons in an atom identifies it as an atom of a
particular element. The number of protons in an atom is referred to as the element’s
atomic number. Remember that because all atoms are neutral, the number of protons and
electrons in an atom must be equal. So, once you know the atomic number of an element,
you know both the number of protons and electrons an atom contains. The average
atomic mass number represents the sum of the atom’s protons and neutrons (both may be
referred to as nucleons). If the atomic number is known you can simply subtract the
number of protons from the average atomic mass number to identify the number of
neutrons that the atom contains.
Potassium
19
K
39.098
Chemical name
Atomic number
Chemical Symbol
Average Atomic Mass
Note: It is incorrect to state that all atoms of a particular element are identical. While it
is true that all atoms of a particular element have the same number of protons and
electrons, the number of neutrons in their nuclei may differ. For example, there are three
different types of potassium one type of potassium atoms contain 20 neutrons, another
contains 21 neutrons and still another contains 22 neutrons; yet all 3 contain 19 protons
and 19 electrons. Atoms such as these are called isotopes. It is important to know that,
even though the numbers of neutrons differ in each isotope, all of these atoms have
essentially the same chemical behavior because chemical behavior is determined by the
number of electrons an atom has.
In nature most elements are found as a mixture of isotopes in a relatively consistent
percentage composition. These isotopes do differ in mass as isotopes containing more
neutrons have a greater mass. The Average Atomic Mass found on the periodic table is
just that, an average of each type of isotopes mass for a particular element in its
percentage composition. To make it easy to identify each of the various isotopes of an
element, chemists add a number after the element’s name. The number that is added is
called the mass number, and it represents the sum of the number of protons and neutrons
in the nucleus for the given isotope (i.e. potassium-19) but the isotope can also be written
using a shortened type of notation involving the chemical symbol, and the atomic
number.
mass number
38
19
atomic number
(this section is Extention Only)
Subatomic particles have extremely small masses that are expressed in scientific notation.
Because scientific notation can be hard to work with, scientists developed a method of
measuring the mass of an atom relative to the mass of a specifically chosen atomic
standard, the carbon-12 atom. Scientists assigned the carbon-12 atom a mass of exactly
12 atomic mass units. Therefore, one atomic mass unit (amu) is defined as 1/12 the
mass of a carbon-12 atom.
The atomic mass of an element is the weighted average mass of the isotopes of that
element. For example, the atomic mass of chlorine is 35.453 amu. Chlorine exists
naturally as a mixture of about 75% chlorine-35 and 25% chlorine-37. Because the
atomic mass is a weighted average, the chlorine-35 atoms, which exist in greater
abundance that the chlorine-37 atoms, have a greater effect in determining the atomic
mass. The atomic mass of chlorine is calculated by summing the products of each
isotope’s percentage abundance times its atomic mass. You can calculate the atomic
mass of any element if you know its number of naturally occurring isotopes, their masses,
and their percent abundances.
i.e. Calculating the Weighted Average Atomic Mass of Chlorine
Clorine-35
atomic mass: 34.969 amu
Percent abundance: 75.770%
Mass contribution:
(34.969 amu)(75.770%)
= 26.496 amu
Clorine-37
atomic mass: 36.966 amu
% abundance: 24.230%
mass contribution:
(36.966 amu)(24.230%)
= 8.957 amu
Weighted average atomic mass of chlorine:
= (26.496 amu + 8.957 amu)
= 35.453 amu
Analyzing an element’s mass can generally give you insight into what the most abundant isotope
for the element may be. If the mass number is a low decimal it suggests most of the isotopes are
in the elements lower form (i.e. 40.078), while a mass number with a higher decimal suggests
that most of the isotopes are in the elements higher form (i.e. 22.990)..
What’s the difference between Mass and Weight? (this section is fair game)
People often mistakenly use the terms mass and weight interchangeably. So how exactly are
they different? We know that matter is anything that has mass and takes up space. Well, mass is
a measurement that reflects the amount of matter an item has. Weight, however, is a
measurement not only of the amount of matter but also of the effect of Earth’s gravitational pull
on that matter. Doesn’t sound that different? Did you know that the earth’s gravitational force is
not exactly the same everywhere on earth, and that it actually becomes less as you move away
from Earth’s surface at sea level? Yo0u might wonder why it is so important to think of matter
in terms of mass. Well the answer is one of practicality; scientists need to be able to compare the
measurements that they make in different parts of the world without the cumbersome task of
calculating varying gravitational forces around the world.
Niels Bohr looked at the quantized energies produced when a hydrogen atom in an
excited state loses energy. From these results, he proposed that the electron
surrounding a hydrogen atom moves around the nucleus in circular orbits. When a hydrogen
atom absorbed energy, the electron jumped from an orbit nearer the nucleus to an orbit farther
away from the nucleus. Likewise, when a hydrogen atom released energy, the electron jumped
from an obit farther away from the nucleus to one closer to the nucleus, releasing a photon on of
light of characteristic color. Although Bohr’s model explained way only light of certain colors is
emitted when hydrogen atoms jump from an excited state to the ground state, it is basically
incorrect. Electrons do not move in circular orbits.
By de Broglie & Schrödinger
(this section is Extension Only)
In the wave mechanical model, electrons are found in locations outside the nucleus called
orbitals. Orbitals are not the same as circular orbits. Unlike our knowledge of planets in orbits,
we cannot know at any one time exactly where an electron is. We can only know where an
electron is likely to be. An electron is likely to be somewhere inside the volume of space
described by the orbital. Chemists have decided to describe the shape of an orbital based on
ninety percent probability. This means that 90 percent of the time, the electron will be found
inside the volume of the orbital, and the other 10 percent, the electron will be found somewhere
outside the orbital volume.
Principle components of the Wave Mechanical Model of the Atom
1. Atoms have a series of energy levels called principal energy levels, which are
designated by whole numbers symbolized by n; n can equal 1, 2, 3, 4, . . . Level 1
corresponds to n=1, level 2 corresponds to n=2, and so on.
2. The energy increases as the value of n increases.
3. Each principal energy level contains one or more types of orbitals, called sublevels.
4. The number of sublevels present in a given principal energy level equals n. For
example, level 1 contains one sublevel (1s); level 2 contains two sublevels (two types
of orbitals – the 2s orbital and the three 2p orbitals), and so on. These are
summarized in the following table.
n
Sublevels (Types of Orbitals) Present
1
1s (1)
2
2s (1) 2p (3)
3
3s (1) 3p (3) 3d (5)
4
4s (1) 4p (3) 4d (5) 4f (7)
The number of each type of orbital is shown in parentheses.
5. The n value is always used to label the orbitals of a given principal level and is
followed by a letter that indicates the type (shape) of the orbital. For example, the
designation 3p means an orbital in level 3 that has two lobes ( a p orbital always has
two lobes)
6. An orbital can be empty or it can contain one or two electrons, but never more than
two. If two electrons occupy the same orbital, they must have opposite spins.
7. The shape of an orbital does not indicate the details of electron movement. It
indicates the probability distribution for an electron residing in that orbital.
Orbital shapes
(a)
(b)
Important terms
Ground state – an atom in its lowest energy state.
Excited state – the state of an electron in an atom with excess energy.
Orbital – A space described by the probability of finding an electron with that particular volume.
1s – the orbital which is closest to the nucleus and of the lowest energy.
Principle energy levels – the major energy levels found in atoms (n=1, n=2, n=3, and so on)
Sublevels – sub divisions of the principal energy levels.
Core electrons – all electrons which are not in the highest energy level (closer to the nucleus).
Valence electrons – electrons in the principal energy level furthest from the nucleus (the highest energy level).
LEWIS STRUCTURES OF THE ELEMENTS (this is fair game)
All of the Group I elements and hydrogen have one and only one electron in the outside shell.
That single electron is what gives these elements the distinctive character of the group. The
Lewis structures are just an attempt to show these valence electrons in a graphic manner as they
are used to combine with other elements. The element symbol is in the center and as many as
four groups of two electrons are shown as dots above, below, to the right and left of the element
symbol to show the valence electrons. All of the inert gases (noble gases) have all eight of the
electrons around the element symbol, except for helium, which has only two electrons even with
a full shell. Below is a demonstration of the noble gases written in Lewis structure. Notice the
electrons are represented as dots around the element symbol.
All the other elements have less than eight electrons in the outside shell. These electrons can be
in the positions of the eight electrons of the noble gases, but there are some suggestions about
where they belong. The Group I elements have only one electron in the outer shell, so it really
does not matter where the electron dot is placed, over, under, right or left of the element symbol.
Group II elements have two electrons. Some authors will place the two electron dots together on
any side of the element symbol because the electrons really are in an s subshell together.
Some authors will show the electrons separated from each other in any of the two
positions with only one electron in each position. The reasoning behind that is that the
electrons really do try to move as far away from each other as possible.
Boron and the elements below it on the periodic table all have three electrons in the
outside shell. These electrons may be grouped as each electron alone in one of the
positions around the element symbol or as a group of two (s) electrons in one position
and one electron in another. Boron is usually shown with separate electrons because it
bonds mostly covalently. Covalent bonds, we know from the shape of molecules, tend to
blend the s and p subshells into sp orbitals with one s and one p orbital blended, sp2
orbitals with one s and two p orbitals blended, or sp3 orbitals, using the single s orbital
with all three p orbitals. The sp2 orbitals of boron tend to be flat trigonal shape, that is,
the bonds are at 120 degrees from each other in a flat circle around the boron atom in the
center. The Lewis structure of boron is any of the shapes below.
Carbon and the elements below it have four electrons in the outer shell. Carbon and
silicon are usually shown in Lewis structures to have four separated electrons, again
because these elements bond purely with covalent bonds. The sp3 orbitals of carbon and
silicon are tetrahedral in shape.
Nitrogen and the elements below it have five electrons in the valence shell, so they must
be shown with one pair (anywhere) and three solitary electrons.
Oxygen and the elements below it have six valence electrons and so must have two pairs
and two solitary electrons.
Elements in the halogen group, Group VII, all have seven electrons in the outer shell, so
only there are three groups of two and a single electron in the last position.
The transition elements and the Lanthanide and Actinide series elements are not often
used in the covalent bonds that the Lewis structures usually portray, but these metal
elements can be portrayed in this manner using the number of electrons in the outer shell
that corresponds with the valence of the element.
In using the Lewis structures to show covalent bonds, the pair of electrons that are in the
bond are shown as a dashed line. For example, ammonia would be shown with the bonds
from the nitrogen to the hydrogens and the unshared pair of electrons on the nitrogen.
Notice that the electrons from all of the participants in this molecule are all accounted for.