Unit+4

= Atomic Structure =

All the matter in the universe is composed of basic substances called elements. These elements can be further divided into atoms. An atom consists of a small, dense nucleus containing all of its protons and neutrons, surrounded by electrons that fill the remaining volume of the atom. An atom is the smallest particle of any element that has the properties of that element.

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All atoms share the same basic structure. Over the past 2000 years, scientists have proposed different models of the atom.

Democritus (in 440 BC) proposed that you would eventually end up with an uncutable piece of matter. He called the piece an atom (from the Greek word //atomos// which means indivisible.)

In 1803, John Dalton of England introduced the atomic idea to chemistry (and is called the Father of Modern Atomic Theory for his efforts). Dalton performed experiments with gases that convinced him that matter was made up of tiny, indivisible particles. He observed, for example, that the same amounts of hydrogen and oxygen always combined to form a given amount of water. He reasoned that each element must be made of its own unique kind of particle an that these particles combine in simple ways. Dalton called these basic particles atoms and pictured then as tiny, solid spheres. Based on his experiments, Dalton developed a theory of the structure of matter. His theory contained four main concepts. In 1897, J.J. Thomson discovered the electron, the first subatomic particle. He also was the first to attempt to incorporate the electron into a structure for the atom. The internal structure of the atom had been a source of speculation for thousands of years.
 * 1. All matter is composed of tiny, indivisible particles called atoms.
 * 2. Atoms of each element are exactly alike.
 * 3. Atoms of different elements have different masses.
 * 4. Atoms of different elements can join together to form compounds.

Thomson faced two major problems: (1) how to account for the mass of the atom when the electron was only about 1/1000 the mass of the hydrogen atom (the more modern figure is 1/1836) and (2) how to create a neutral atom when the only particle available was negatively charged. He believed that the atom had to contain positively charged particles to balance the negative charge of the electrons. But in all his experiments, Thomson was never able to find these positively charged particles. So he proposed a model of the atom that is sometimes referred to as Thomson's "plum pudding model." Each atom was pictured as being made of a pudding-like, positively charged material. Scattered in this material, like plums in a pudding, were the negatively charged electrons. the electrons. Further experiments led Ernest Rutherford to a different view. In 1911, Rutherford bombarded a very thin piece of gold foil with a stream of positively charged alpha particles. Alpha particles are helium atoms without any electrons. He reasoned that if atoms were solid, the bombarding particles should bounce right off. He found, to his surprise, that most of the alpha particles went straight thorough the gold foil as if there was nothing there. A few alpha particles bounced back very strongly. From the fact that most of the alpha particles passed through the foil, he concluded that an atom is mostly empty space. From the fact that some of the alpha particles bounced back, he concluded that the center of an atom has a heavy positively charged core, which is called the nucleus. In 1913, Niels Bohr proposed a model that described the arrangement of electrons in atoms. In the Bohr planetary model, the electrons move in orbitals at different distances from the nucleus. In this way the electrons are like planets which move around the sun. This model was very successful in explaining many of the characteristics of the hydrogen atom. However, it was difficult to apply this model to ore complex atoms. In place of the sun was small ball of positive charge. Around this moved negative particles in much the same way that planets orbit the sun.

Present models of the atom show electrons that move very quickly about the nucleus but not in circular orbits. The region around the nucleus occupied by electrons is an electron cloud. Scientists have calculated the most likely locations of an electron within the electron could. The density of the could decreased farther from the nucleus, because electrons are less likely to be found there. They are more likely to be found close to the nucleus. the electrons making up the could should be thought of as being everywhere in the cloud at once. this is much like the whirling blades of a fan. The blades seem to fill the space around the center of the fan. However we know the blades are not everywhere at the same instant. Not all the electrons in an electron cloud have the same energy. The farther an electron is from the nucleus, the more energy the electron has. Each electron is in an energy level. Electrons with the lowest amount of energy are in the first energy level. This energy level is closest to the nucleus. There is a limit on how many electrons can be in each energy level. The lowest energy level can never hold more than 2 electrons. The atom's second energy level can have up to 8 electrons. The third energy level can hold up to 18 electrons. Energy levels farther from the nucleus can have 32 or more electrons.

There are two more restrictions on the number of electrons in energy levels. First, the number of electrons in a neutral atom is the same as the number of protons in its nucleus. This limits the total number of electrons. Second, there are no more than eight electrons in the highest occupied energy level. Energy levels without electrons are not counted.

Subatomic Particles
An **electron** is a particle that moves around the nucleus forming a cloud of negative charge. A **proton** is a particle that gives the nucleus its positive charge. A **neutron** is a particle with no charge. Neutrons are also in the nucleus of the atom. A proton and neutron are equal in mass. However, both have masses more than 1800 times that of an electron. All atoms of an element have the same number of protons. For example, every hydrogen atom has one proton in its nucleus. Every carbon atom has six protons. The number of protons in the nucleus of an atom determines what the element is. The number of protons in an atom is called the **atomic number**.

Build your own atoms here!

Atomic Mass
The mass of an atom depends on the number of protons and neutrons in its nucleus. Compared to the nucleus of an atom, the electrons have very little mass. The mass of a single proton or neutron itself is extremely small. A proton has a mass of 0.000 000 000 000 000 000 000 001 673 grams. Scientists have agreed upon a standard atomic mass unit to measure the mass of atoms. Protons and neutrons have a mass of about 1 amu each. Electrons are much smaller and have a mass of 1/1800 amus. The mass number of an atom is simply the sum of the number of protons and neutrons. If you know the mass number and the atomic number, you can find the number of neutrons. The number of neutrons is found by subtracting the atomic number from the mass number.

Isotopes
All atoms of an element have the same number of protons. However, some of these atoms may have more or fewer neutrons than others. Atoms of the same element with different numbers of neutrons are called isotopes. There are two ways to show the difference between isotopes of an element. In one way, the name of the element is followed by the mass number. For example, oxygen-16 stands for an atom of oxygen that has a mass number of 16. The second way is to write the symbol of the element with the mass number and atomic number. For example is agron-39. Most elements in nature are found to be mixtures of isotopes.
 * 39 || Ar ||
 * 18 ||^  ||

= Symbols = = = Atoms are assigned the first letter of their name as a convenient way of referring to them. The capital letter is used, but if a second or third letter is necessary, then lower case letters are used.

Some elements have been given symbols based on their names in Latin, which was once commonly used by scientists as an international language. For example, here are some familiar elements with their Latin names:


 * ~ Element ||~ Symbol ||~ Latin name ||
 * sodium || Na || //natrium// ||
 * potassium || K || //kalium// ||
 * iron || Fe || //ferrum// ||
 * copper || Cu || //cuprum// ||
 * silver || Ag || //argentum// ||
 * tin || Sn || //stannum// ||
 * gold || Au || //aurum// ||
 * mercury || Hg || //hydrargyrum// ||
 * lead || Pb || //plumbum// ||

The Bohr model is a planetary model of the atom. Neils Bohr proposed that the electrons orbiting the atom could only occupy certain orbits. In the Bohr model, the most stable, lowest energy level is found in the innermost orbit. This first orbital forms a shell around the nucleus and is assigned a principal **quantum number** (n) of n=1. Additional orbital shells are assigned values n=2, n=3, n=4, etc. Increasing numbers of electrons can fit into these orbital shells according to the formula 2n 2. The first shell can hold up to two electrons, the second shell (n=2) up to eight electrons, and the third shell (n=3) up to 18 electrons. Subshells or suborbitals (designated //s, p, d,// and //f//) with differing shapes and orientations allow each element a unique electron configuration. Bohr's work earned a Nobel Prize in 1922. Subsequently, more mathematically complex models based on the work of French physicist Louis Victor de Broglie (1892-1987) and Austrian physicist Erwin Schrodinger (1887-1961) that depicted the particle and wave nature of electrons proved more useful to describe atoms with more than one electron. The standard model incorporating quark particles further refines the Bohr model. Regardless, Bohr's model remains fundamental to the study of chemistry, especially the **valence** shell concept used to predict an element's reactive properties.
 * || **Energy Level**
 * (Principal Quantum Number) ** ||  || **Shell Letter** ||   || **Electron Capacity** ||
 * 1 ||  || K ||   || 2 ||
 * 2 ||  || L ||   || 8 ||
 * 3 ||  || M ||   || 18 ||
 * 4 ||  || N ||   || 32 ||
 * 5 ||  || O ||   || 50 ||
 * 6 ||  || P ||   || 72 ||
 * 3 ||  || M ||   || 18 ||
 * 4 ||  || N ||   || 32 ||
 * 5 ||  || O ||   || 50 ||
 * 6 ||  || P ||   || 72 ||
 * 5 ||  || O ||   || 50 ||
 * 6 ||  || P ||   || 72 ||
 * 6 ||  || P ||   || 72 ||
 * 6 ||  || P ||   || 72 ||