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Atomic Theories The idea of an atom -- the smallest particle of matter -- has intrigued mankind since the beginning of civilization. Throughout the centuries the "view" of the atom has changed. New ideas, and new technologies have influenced the model of the atom. This view of the atom is still a Theory and therefore it is still subject to change. The modern model of the atom is called the Quantum Model and you will study this model in future. The chart below summarizes the various atomic models that have been developed during the course of history. Scientist & approximate Date Democritus c.300 BC Name of Model, Sketch and main idea of theory Atom the indivisible particle Atomos (in ancient Greek) means "that which cannot be further broken down into smaller pieces". The solid sphere model Atoms are seen as solid, indestructible spheres (like billiard balls) Dalton c.1800 Importance and Improvement on previous model Talks about the atom as the smallest particle of matter. Defines the atom as an indivisible particle Explains certain natural occurrences such as the existence of elements Explains a lot of chemical properties such as how atoms combine to form molecules. Explains chemical change better than the Particle Theory Confirms the basic Laws of Chemistry: Conservation of Mass & definite Proportions Shortcomings Problems or why was it changed Does not give a scientific view of the atom only a conceptual definition Does not talk about subatomic particles (Electrons, Protons, Neutrons) Does not include the existence of the nucleus Does not explain the existence of ions or isotopes Does not talk about subatomic particles (Electrons, Protons, Neutrons) 1 The raisin bun Model or the chocolate chip cookie model: Atoms are solid spheres made-up of a solid positive mass (or core) with tiny negative particles embedded in the positive core. Infers on the existence of electrons and protons. Introduces the concept of the nucleus. Ifers on the relative nuclear density and atom mass of different atoms Does not explain the existence of electrons outside the nucleus does not explain the role of electrons in bonding. Does not talk about neutrons therefore can't explain radioactivity and the existence of isotopes First real modern view of the atom. Explains why the electron spins around the nucleus (Bohr's Contribution) Proposes that the atom Famous Gold Leaf is really mostly empty Experiment proves space that the nucleus is positive and the electrons are outside the nucleus. Does not place electrons in definite energy levels around the nucleus. Doesn't include neutrons in the nucleus. Electrons in Definite energy Levels around the nucleus Used atomic spectra to prove that electrons are placed in definite orbitals (called shells) around the nucleus. It does not explain the shapes of molecules or other abnormalities that result form unevenly shared pairs of electrons (such as the abnormal behaviour of water, the difference in Carbon-Carbon Bonds between diamond and graphite etc..) J.J. Thomson c.1850 The Planetary Model Rutherford c. 1905 (Neils Bohr) BohrRutherford c. 1920 Explains the role of valence electrons in bonding. Relegates the number of valence electrons to the Periods of a periodic table. Fully explains ionic and covalent bonding. Places electrons in definite energy levels 2 e- in the first 8 e- in the second 8 e- in the third (see example below) Does Not relate the valence electrons atomic charge 2 Modern Theory Many Scientists Contributed. Some of the more famous are: Schroedinger Einstein Luis De Broglie Max Planck Frank Hertz Maxwell Fermi Quantum Mechanical Model or Electron Cloud Model The analogy here is that of a "beehive" where the bees are the electrons moving around the nucleus in a "cloud" of energy levels. Advanced Theories will explain bonding and other facts about the behaviour of atoms and their chemical and physical properties in forming new compounds. Other important facts about the particles of an atom: Subatomic Symbol Charge Relative Mass* Location Particle Proton p+ positive 2000 nucleus orbits around Electron enegative 1 nucleus Neutron n0 neutral (zero) 2000 nucleus The mass number and the atomic number from the Periodic Table are very important numbers because they tell us how many subatomic particles are contained in a given atom. The atomic number tells us the number of protons., i.e. Atomic Number = Number of Protons. The atomic mass tells the total number of particles in the nucleus, i.e. Atomic Mass = # of protons + number of neutrons. For example: The square where the element Boron is located on the Periodic Table looks like this From this we can obtain the following information: 3 Element Symbol Boron B Atomic Number 5 Atomic # of # of # of neutrons Mass protons electrons 11 5 5 11 - 5 = 6 Recall that the Bohr-Rutherford Model places the electrons around the nucleus in definite orbitals or energy shells as summarized by the table below. Energy Shell 1 2 3 4 Name of Shell Maximum Number or Symbol Electrons it can contain K 2 L 8 M 8 N 18 of Now we use this and the information from the Periodic Table to draw a Bohr-Rutherford diagram for the Boron Atom as illustrated below: To draw Bohr-Rutherford diagram for Boron we place the first 2 electrons in first shell. The first shell can only hold a maximum of 2 electrons so we start filling the second shell The Boron atom has 5 electrons therefore we have another 3 electrons to place. We place these electrons in the Second shell and we space them apart from one another If there are more than four electrons in the second shell (as in the case of the Fluorine atom), we pair the electron up. This pairing of electrons is explained by more advanced theories which propose that to counteract the repulsive forces between the electrons' negative charges, one elctron spins in the opposite direction of the second electron. Example: Draw a Bohr-Rutherford diagram for the element Sodium. Solution: Using the periodic table we obtain the following information about the sodium atom: Element Symbol Sodium Na Atomic Atomic of Number Mass protons 11 23 11 of electrons 11 of Atomic Diagram neutrons 12 4 Electronic Configuration Electronic Configuration is the method of listing the locations of all the electrons on a specific atom. The system locates each electron by energy level and orbital type. The number of electrons in each orbital type is indicated with a superscript. Orbitals Orbitals are probability diagrams. Specifically, an orbital describes a region in space where there is a 90% change of finding an electron. The electron is never restricted to an orbital as in travels around a nucleus, but it seems to keep returning to this particular region even though its behavior is random. The concept of the orbital differs from Bohr's concept of the orbit. Bohr considered an orbit to be a path that the electron always followed much like a train stays on a track. The concept of the orbital was developed in Schrodinger's work to avoid violating the Heisenberg Uncertainty Principle. In the Modern Theory of Atomic Structures a picture of an orbital is also called a Probability Diagram. By agreement among chemists, the orbital is a 90% Probability Diagram. This idea allows the electron to be found anywhere and still indicates where the electron spends most of its time. Their names, and shapes, are based on the value of the Azimuthal Quantum Number. l = 0, the orbital is called "s". l = 1, the orbital is called "p". l = 2, the orbital is called "d". l = 3, the orbital is called "f". Filling the electron energy levels Main Energy Number of electrons in level in each sub-level This is a useful pattern of remembering how the levels are filled. http://www.webchem.net/notes/atomic_structure/electron_configuration.htm http://www.lenntech.com/periodic-chart.htm http://acswebcontent.acs.org/periodic/tools/PT.html http://www.webelements.com/ 5 PERIODIC PROPERTIES Description: The properties of the elements exhibit trends and these trends can be predicted with the help of the periodic table. They can also be explained and understood by analyzing the electron configurations of the elements. This is because, elements tend to gain or lose valence electrons to achieve the stable octet formation. In addition to this activity, there are two other important trends. First, electrons are added, one at a time, moving from left to right across a period. And, as this happens, the electrons of the outermost shell experience increasingly strong nuclear attraction. As a result, the electrons become closer to the nucleus and more tightly bound. The second trend is the moving down a column in the periodic table, where the outermost electrons become less tightly bound to the nucleus. And these trends explain the periodicity observed in the elemental properties of atomic radius, ionization energy, electron affinity, and electronegativity. But, before going into that we need to know a bit more about the above mentioned terms: Atomic Radius The atomic radius of an element is half of the distance between the centers of two atoms of an element that are in contact with each other. Generally, the atomic radius decreases across a period, from left to right and increases down a given group. Therefore, the atoms with the largest atomic radii are located in Group I and at the lower half of groups. . . Ionization Energy Ionization energy or ionization potential is the energy required to completely remove an electron from a gaseous atom or ion. And, the closer and more tightly an electron is bound to the nucleus, the more difficult it is to remove and the higher its ionization energy. Ionization energy is also required to remove a second valence electron from the univalent ion to form the divalent ion, and so on. A. Increase across the period (L R) due to: i) ii) More protons Decreasing radius stronger attraction B. Decrease IE down the group due to: More electrons shielding less attraction thus lower IE Electron Affinity Electron affinity is the energy change that occurs when an electron is added to a gaseous atom. It reflects the ability of an atom to accept an electron. And the atoms with stronger effective nuclear charge have a greater electron affinity. Therefore, some generalizations can be made about the electron affinities of certain groups in the periodic table. The alkaline earths have low electron affinity values. This is because they have filled sub shells. But, the halogens have high electron affinities because of the 6 addition of an electron to an atom results in a completely filled shell. Noble gases have zero electron affinities, since each atom possesses a stable octet and will not accept an electron readily. Electronegativity An atom with higher electro negativity has a great capacity for attracting bonding electrons. Therefore, electronegativity is a measure of the attraction of an atom for the electrons in a chemical bond. Trends for Electronegativity:Electronegativity increases across the period. It decreases down the group. Reasons for trends:Increase across periods due to more protons, stronger attraction, small radius thus large electronegativity. Decrease down the group due to increase in the number of electrons an atom already has, which means the atom is less able to attract another atom’s electrons, thus small electronegativity. periodic trends at form d n u po te r com arac t h n c e l c l li Cova meta ency electron affinity end ion t Ionization energy electron affinity electronegativity atomic and ionic radii Ionization energy electronegativity atomic and ionic radii . CHEMICAL BOND The process of two or more atoms joining together to form a molecule is called bonding. In general, bonding is a chemical change that occurs during chemical reactions. During bonding electrons are involved in "fusing or gluing" two or more atoms together. There are three general ways in which electrons play a role in bonding. They can either be electrostatically moved from the atom of one element to another atom of a second element. This is known as ionic bonding They can be shared between two different atoms or they can be shared among the same element of the same atom. This sharing type of bonding is called molecular or covalent. Some elements are surrounded by a cloud of free electrons which are shared among all atoms of the same element. This mainly takes place in metallic elements and explains the many properties of metals. This type of bonding is called metallic bonding. 7 The physical behavior of materials is thought to be determined by the structure of the material. The phrase "Structure Determines Function/Behavior" is often used as a reminder that the observable behavior of a material depends on its structure. Metals The metallic elements are found on the left-hand side of the Periodic Table. Metals are lustrous, malleable, and ductile. They conduct heat and electricity quite well. Metals tend to have high melting points with several notable exceptions of which mercury(Hg) is the most well known. Metals are not soluble in water. The proposed structure for samples of metals is that of the "band theory" of metallic bonding. Metal samples are considered to be composed of metal atoms that are held together by delocalized bonds formed by all of the atoms in the sample. It is proposed that electrons can move easily within the sample. An earlier theory was sometimes called the "electron sea model" which pictured the sharing of valence electrons by all atoms in the sample. Properties of Metals: Solid and liquid samples of metals can conduct an electric current due to the mobility of the valence electrons of the atoms in the sample. Most metals will have high melting points. Metals are lustrous(shiny), malleable(able to flattened into sheets) and ductile(able to be drawn into wires). Ionic Compounds Ionic compounds are formed when metallic elements from the left-hand side of the Periodic Table react with nonmetallic elements from the right-hand side of the Table. Ionic compounds have high melting and boiling points. Many ionic compounds are soluble in water. Ionic compounds are brittle and do not conduct electricity when in the solid form. The molten (liquid) form of ionic compounds will conduct an electric current. In a molten sample of an ionic compound the ions are free to move. The cations can move to the negative electrode and acquire electrons. Similarly, the anions can move to the positive electrode and release electrons, thereby causing a flow of electricity in the outer circuit. The proposed structure for ionic compounds is that positive cations ( formed from the elements on the left-hand side of the Periodic Table) attract negative anions(formed from the elements on the right-hand side of the Periodic Table) to form a network or lattice of oppositely charged ions arranged in three-dimensional patterns that depend on the size and charge of the ions. Properties of Ionic Compounds Solid ionic compounds do not conduct an electric current. Molten samples of ionic compounds can conduct an electric current due to the mobility of the ions which are free to move to the electrodes and react. Ionic compounds have high melting and boiling points. Many ionic compounds can dissolve in water. Dissolved ionic compounds separate into cations and anions in solution. The mobile ions move to the electrodes and react to accept and release electrons creating a flow of electricity in the outer circuit. Ionic compounds that are water soluble are strong electrolytes. Covalent Compounds Molecular compounds form when two or more nonmetal atoms form units that are called molecules. Molecular compounds generally have low melting and boiling points. 8 Molecular compounds do not conduct electricity in the solid form or in the liquid form. Some molecular compounds dissolve in water and some do not. The proposed model for molecular compounds is that atoms of nonmetallic elements can bond by sharing electrons to form units called molecules. A very common example is the water molecule that forms when an oxygen atom shares electrons with two hydrogen atoms to form one unit that is called a water molecule. A sample of water is thought to be a set of these molecular units. Molecules are attracted to other molecules by attractions that are called Intermolecular Attractions (attractions between molecules). In a physical change such as melting or evaporation these intermolecular attractions must be broken. Each molecule is still a molecular unit but it is separated from the other molecular units. Intermolecular attractions are generally weaker attractions and do not require as much energy to break as do metallic, ionic or covalent bonds. Therefore, molecular compounds generally have low melting points and low boiling points. Molecules may be polar molecules which have a charge separation due to the shape of the molecule and the polarity of the bonds between the atoms forming the molecule. These molecules are pictured as having a positive end and a negative end and are said to be dipoles or polar molecules. Other molecules are non-polar molecules. These are molecules that have no charge separation again due to the shape of the molecule and the polarity of the bonds between the atoms forming the molecule. These non-polar molecules are pictured as having no dipole character. LIKE DISSOLVES LIKE. Polar molecules tend to dissolve materials with charged nature such as ionic compounds and other polar molecules. Likewise, non-polar molecules tend to dissolve other non-polar molecules. Properties of Molecular Compounds Solid molecular compounds do not conduct an electric current. Molten samples of molecular compounds do NOT conduct electricity. Molecular compounds have low melting and boiling points. Polar molecular compounds can dissolve in water. Some very polar molecular compounds can be "ripped" into ions by the water molecules. These are said to be "dissociated" into ions by the dipole nature of the water molecules. Polar molecular compounds are usually not soluble in non-polar solvents. Non-polar molecular compounds are usually not soluble in water. They are soluble in other non-polar solvents such as toluene and other organic solvents. Polar molecular compounds that are dissociate into ions by water electrolytes. http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/bom1s2_11.swf http://www.chemguide.co.uk/index.html#top http://www.ktf-split.hr/periodni/en/abc/e-config.html#7Period 9