ionic compound.
An ionic compound contains positively and negatively charged ions
It should be pointed out that the Na+ and Cl- ions are not chemically bonded together. Whereas atoms in molecular compounds, such as H2O, are chemically bonded.
Ionic compounds are generally combinations of metals and non-metals.
Molecular compounds are general combinations of non-metals only.
Pure ionic compounds typically have their atoms in an organized three dimensional arrangement (a crystal). Therefore, we cannot describe them using molecular formulas. We can describe them usingempirical formulas.
If we know the charges of the ions comprising an ionic compound, then we can determine the empirical formula. The key is knowing that ionic compounds are always electrically neutral overall.
Therefore, the concentration of ions in an ionic compound are such that the overall charge is neutral.
In the NaCl example, there will be one positively charged Na+ ion for each negatively charged Cl- ion.
What about the ionic compound involving Barium ion (Ba2+) and the Chlorine ion (Cl-)?
1 (Ba2+) + 2 (Cl-) = neutral charge
Resulting empirical formula: BaCl2 [2]
Atomic structure
The Nucleus
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The atomic nucleus is the central area of the atom. It is composed of two kinds of subatomic particles: protons and neutrons.
Diagram showing the atomic structure with the protons and neutrons held together to form the dense area of the nucleus.
Atoms are the building blocks of all matter. Everything you can see, feel and touch is all made of atoms. There are even things you cannot see, feel, hear or touch that are also made of atoms. Basically, everything is made up of atoms.
In 1909, Ernest Rutherford led Hans Geiger and Ernest Marsden through what is known as the Gold Foil Experiments. During the experiments they would shoot particles through extremely thin sheets of gold foil. In 1911, Rutherford came to the conclusion that the atom had a dense nucleus because most of the particles shot straight through, but some of the particles were deflected due to the dense nucleus of the gold atoms. This theory would eliminate the idea that the atom was structured more like plum pudding. The plum pudding model was the leading model of atomic structure until Rutherford's findings.
Atomic Numbers
The atomic nucleus is in the center of the atom. The number of protons and neutrons in the atom define what type of atom or element it is. An element is a bunch of atoms that all have the same type of atomic structure. For instance, hydrogen is an element. Every hydrogen atom is made up of 1 proton, 0 neutrons, and 1 electron.
The composition of the atomic nucleus gives us lots of information about the element it represents. The number of protons inside the nucleus gives us theatomic number. The protons have a positive (+) charge. In order for the atom to have a neutral charge, the electrons (-) need to balance it out with their negative charge. Therefore, in a neutral atomthere are just as many protons as electrons. So, if you know the atomic number and know the charge of the atom then the number of electrons is easy to find. For instance, hydrogen has 1 proton, 1+, so in order for the hydrogen atom to be neutral it must have 1- charge. Therefore, hydrogen has 1
electron.
Where do the neutrons fit in all of this? Well,neutrons are neutral. To keep it all straight I use the first letters: Neutrons are Neutral, and Protons arePositive. I then remember Electrons through the process of Elimination.
Although the neutrons do not give the atom any charge, they still hold their own weight in the importance of the atomic structure. The neutron is the largest of the subatomic particles. When put the neutrons and protons together we get the atomic mass . The electrons are so small that their mass only counts for .01%. The electrons are not inside of the nucleus; instead they are flying around like crazy on the outside of the nucleus.
Since the atomic number gives us the number of protons in an atom and the atomic mass gives us the number of protons and neutrons, we can find the number of neutrons by subtracting the atomic number from the atomic mass.
Atomic mass - atomic number = number of neutrons.
The atomic number of an atom gives each element its identity. You can find out which element it is by its atomic number and reverse the process to find out what the atomic number is if you know which element you are working with.
Let's run through all of the numbers with an element, oxygen.
Oxygen
Atomic Number: 8
Atomic Mass: 16 [3]
The ability of atoms to lose or to gain electrons.
Next, let's review two atomic properties important to bonding that are related to the position of the element on the periodic table. They are the tendency or ability of atoms tolose electrons and the tendency or ability to gain electrons.
First, let's consider the ability to lose electrons. This is related to ionization energy, which you studied in a previous lesson. The ionization energy, of course, is the amount of energy that it takes to remove an electron from an atom. You have learned that the ionization energies are lowest for the elements down and on the left hand side of the periodic table and increase as you go up and all the way across to the right including the inert gases.
The ionization energy measures how hard it is to lose or remove an electron. High ionization energy means that it is hard to lose electrons. Low ionization energy means that it easy to lose electrons. The elements on the left side lose their electrons fairly easily and the elements on the right side of the periodic table do not lose their electrons very easily. Taking vertical position on the table into account, the elements that are lower on the table lose electrons more easily and the elements that are higher have a harder time losing electrons. Thus the overall trend is from most easily losing electrons on the lower left to least easily losing electrons on the upper right. Keep that trend in mind.
Ability to Lose Electrons
The ability to gain electrons is also related to the position on the periodic table. You should recall that as you go from left to right on the periodic table, the attraction for electrons increases and the ability to gain electrons increases. This is true all the way across the periodic table except/em> for the inert gases. There is an abrupt drop in the ability to gain electrons when we get to the inert gases. This is because their energy level is full and any additional electrons will have to start a new energy level.
Ability to Gain Electrons
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