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Although atoms are the smallest unique unit of a particular element, in nature only the noble gases can be found as isolated atoms. Most matter is in the form of ions, or compounds.

Molecules and chemical formulas

A molecule is comprised of two or more chemically bonded atoms. The atoms may be of the same type of element, or they may be different.

Many elements are found in nature in molecular form - two or more atoms (of the same type of element) are bonded together. Oxygen, for example, is most commonly found in its molecular form "O₂" (two oxygen atoms chemically bonded together).

Oxygen can also exist in another molecular form where three atoms are chemically bonded. O₃ is also known as ozone. Although O₂ and O₃ are both compounds of oxygen, they are quite different in their chemical and physical properties. There are seven elements which commonly occur as diatomic molecules. These include H, N, O, F, Cl, Br, I.

An example of a commonly occurring compound that is composed of two different types of atoms is pure water, or "H₂O". The chemical formula for water illustrates the method of describing such compounds in atomic terms: there are two atoms of hydrogen and one atom of oxygen (the "1" subscript is omitted) in the compound known as "water". There is another compound of Hydrogen and Oxygen with the chemical formula H₂O₂ , also known as hydrogen peroxide. Again, although both compounds are composed of the same types of atoms, they are chemically quite different: hydrogen peroxide is quite reactive and has been used as a rocket fuel (it powered Evil Kenievel part way over the Snake River canyon).

Most molecular compounds (i.e. involving chemical bonds) contain only non-metallic elements.

Molecular, Empirical, and Structural Formulas

Empirical vs. Molecular formulas

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  Molecular formulas refer to the actual number of the different atoms which comprise a single molecule of a compound.
  
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  Empirical formulas refer to the smallest whole number ratios of atoms in a particular compound.
  

Sometimes the molecular formulas are drawn out as structural formulas to give some idea of the actual chemical bonds which unite the atoms.

Structural formulas give an idea about the connections between atoms, but they don't necessarily give information about the actual geometry of such bonds.

The nucleus of an atom (containing protons and neutrons) remains unchanged after ordinary chemical reactions, but atoms can readily gain or lose electrons.

For example, Sodium (Na) has 11 protons and 11 electrons. However, it can easily lose 1 electron. The resulting cation has 11 protons and 10 electrons, for an overall net charge of 1+ (the units are electron charge). The ionic state of an atom or compound is represented by a superscript to the right of the chemical formula: Na⁺, Mg²⁺ (note the in the case of 1+, or 1-, the '1'is omitted). In contrast to the Na atom, the Chlorine atom (Cl) easily gains 1 electron to yield the chloride ion Cl^- (i.e. 17 protons and 18 electrons).

Na⁺ and Cl^- are simple ions, in contrast to polyatomic ions such as NO₃^- (nitrate ion) and SO₄^2- (sulfate ion). These are compounds made up of chemically bonded atoms, but have a net positive or negative charge.

The chemical properties of an ion are greatly different from those of the atom from which it was derived.

Predicting ionic charges

Many atoms gain or lose electrons such that they end up with the same number of electrons as the noble gas closest to them in the periodic table.

The noble gasses are generally chemically non-reactive, they would appear to have a stable arrangement of electrons.

Nitrogen has an atomic number of 7; the neutral Nitrogen atom has 7 protons and 7 electrons. If Nitrogen gained three electrons it would have 10 electrons, like the Noble gas Neon (10 protons, 10 electrons). However, unlike Neon, the resulting Nitrogen ion would have a net charge of N^3- (7 protons, 10 electrons).

This is mainly true for the elements on either side of the chart.

Ions form when one or more electrons transfer from one neutral atom to another. For example, when elemental sodium is allowed to react with elemental chlorine an electron transfers from a neutral sodium to a neutral chlorine. The result is a sodium ion (Na⁺) and a chlorine ion, chloride (Cl^-):

The oppositely charged ions attract one another and bind together to form NaCl (sodium chloride) an ionic compound.

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.

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.

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

Diagram showing the atomic structure with the protons and neutrons held together to form the dense area of the nucleus.

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.

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.

Types of Chemical Bonds

A group of atoms bonded to one another form a molecule.

If the molecule has more than one type of element present it is a compound.