Chemistry 103A; Sections 5, 6, 7, 8; Lecture 4; 28 Aug 00

The structure of atoms, continued

What's inside an atom? Or what are the components of an atom?

Atoms are composed of elementary particles:

protons - "heavy," positive electric charge (mass = 1.6726 ´ 0-24g)
neutrons - "heavy," electrically neutral (mass = 1.6749 ´ 10-24g)
electrons - "light," negative electrical charge (mass = 9.1094 ´ 10-28g)
It is easier to measure relative masses that to measure absolute masses so we introduce the atomic mass unit (amu) as 1/12 the mass of the carbon atom which has 6 protons and 6 neutrons in its nucleus. In amu the masses of the elementary particles are:
protons – 1.007276 amu
neutrons – 1.008665 amu
electrons – 0.0005485799 amu
Electron and proton charges are equal and opposite - they "balance" each other.
Neutrons have zero charge.

In a neutral atom, # electrons = # protons =atomic number

# protons + # neutrons = mass number (or atomic mass number)

mass of proton » mass of neutron » 1836 ´ mass of electron

(Most of the mass of an atom is in the protons and neutrons.)

The protons and neutrons are bound tightly together to form the nucleus. Most of the mass of the atom is concentrated in the nucleus.

The electrons surround the nucleus in a "charge cloud." Since the electrons are 1836 times lighter than protons and neutrons, only about 0.03% of the mass is in the electrons.

On the other hand, the nucleus is very small, most of the size (volume) of the atom is provided by the electrons.

Examples:

Hydrogen: The diameter of the nucleus (proton) is about 0.01 pm, but the diameter of the atom is about 104 pm.

Note: The mass of the electron is 9.1 x 10-28 g while the mass of the H atom is about 1.7 x 10-24 g.

Lead: The diameter of the nucleus is about 0.067 pm, but the diameter of the atom is about 350 pm. (Compare the earth and the UA campus.)

Since the electrons surround the nucleus (they are on the "outside,") bonding must depend on what the electrons are doing.

(The electrons "shield" the nucleus – the electrons "bump into" stuff before the nucleus does.)

(There were other models of the atom, e.g. the "raisin muffin" model, but they didn't stand up to experimental test.)

We will look at what the electrons are doing (the electronic structure of the atom) later.
 

Structure of the atomic nucleus

The atomic number

= # protons (always)
= # electrons (in a neutral atom)
determines what an element is.
O is atomic number 8
H is atomic number 1
He is atomic number 2
C is atomic number 6
Cl is atomic number 17
Ar is atomic number 18
etc.


Ions

Sometimes the number of electrons does not equal the number of protons. In this case the atom is no longer electrically neutral. The atom has a charge, it is called an ion.

If it has an excess of electrons it will be a negatively charged ion (an anion). If has a deficiency of electrons it will be a positively charged ion (a cation).

Examples

H+
H-
O2-
Na+
Ca2+
We will find out how and why atoms can form stable ions in Chapter 8.
 

Nuclear structure

Recall:

Mass Number = # neutrons + # protons
Atomic Number = # protons
We write a symbol of the atom which contains all of the nuclear and charge
information as follows:
X is the symbol of the element and the subscript and superscripts tell us about the nuclear structure.
MN = mass number (Sometimes this is called the atomic mass number.)

AN = atomic number

chg = charge

Examples:
1 proton + 0 neutron + 1 electron = 
2 proton + 2 neutron + 1 electron = 
6 proton + 6 neutron + 7 electron = 
82 proton + 125 neutron + 80 electron = 

and so on.

Isotopes

Although atoms of a given element always have the same number of protons, they can have different numbers of neutrons.

(Notice that this will not change the chemical properties, because the number of electrons will be the same.)

Atoms of an element with different mass numbers are called isotopes.

We use the same type symbols to specify the different isotopes:

Examples:

Isotopes of hydrogen - , , 
( is called "deuterium" and  is called "tritium." Not all isotopes have their own special name.)

Isotopes of helium - , 
( is called "helium – 3," and  is called "helium – 4," and so on for the rest of the elements.)

Isotopes of carbon - , , 
(Carbon – 12, carbon – 13, and Carbon 14,etc.)
There are 7 known isotopes of C with mass numbers ranging from 10 to 16

Isotopes of oxygen - , , 
(There are 8 isotopes of oxygen: 13 to 20.)

and so on.

Notice that the symbol of the element and the atomic number subscript are redundant. We leave the subscript on when we are writing nuclear reactions (Chapter 24), but sometimes the atomic number is left off.
For example,  can also be written .
Let’s consider the nuclear structure of .
This symbol tells us that the atom of uranium - 235 contains 92 protons and 235 -  92 = 143 neutrons.  Since there is no charge indicated, we conclude that there are also 92 electrons in the atom.
The Handbook of Chemistry and Physics, sometimes known as The Chemical Rubber Handbook, contains a table of all know isotopes of all known elements. The list runs to about 3,000 isotopes. Not all of the known isotopes are stable, most are radioactive.
 

Why are atomic weights or the atomic masses in the periodic table or in our element lists not integers?

Two reasons:

1. The neutrons and protons in different elements (and even different isotopes of the same element) do not all weigh the same and they do not weigh the same as free neutrons (1.008665 amu) and free protons (1.007276 amu).

These differences are small, but it means that we can’t find the mass of a given isotope by adding the masses of all the neutrons and protons.

2. The main reason is that the atomic weight that we measure is the weighted average of the naturally occurring isotope masses.

Sometimes we call the weighted average of the naturally occuring isotopes the chemical atomic weight.

(With a mass spectrometer chemists can determine the natural abundance of the isotopes of an element and the actual mass of each isotope. The natural abundance is the percentage of each isotope in a sample of the naturally occurring element.)

Examples:

  - 98.89%, mass = 12.0000 amu (by definition)

  - 1.11%, mass = 13.00335 amu

The weighted average is:
Another example:
  - 99.759%, mass = 15.99491 amu
  - 0.037%, mass = 16.995 amu
  - 0.204%, mass = 17.9943 amu
The weighted average is:
One more example:
  - 75.53%, mass = 34.96885 amu
  - 24.47%, mass = 36.9671 amu
Weighted average is:
The weighted averge atomic masses are all in units of amu.
 
 

Overview of the Periodic Table

(We will discuss why the table has this form in more detail in Chapter 8.)

General principles:

Elements in the same column have similar chemical properties.

Reactivity gets more varied as you go down a column.

The left side elements are metals and the right side elements are nonmetals.

Alkali metals
The elements in the far left column are called alkali metals.  They are very reactive metals and tend to form ions with a charge of +1 in compounds.
Alkaline earth metals
The elements in the second from the left column are called alkaline earth metals.  They are not as reactive as the alkali metals and tend to form ions with a charge of +2 in compounds.
Transition elements
The next ten columns comprise the transition elements.  We will learn why they are called that in Chapter 8.  These elements are all metals.
B group

C group

For example, the elements of the boron group all form compounds with hydrogen with the formula, XH4.
N group

O group

Halogens (acid formers)
 
 

noble gases