Atomic+structure+&+bonding

toc =Matter and the Atom=

Dalton’s Atomic Theory
In 1804, John Dalton first proposed the idea that tiny particles called atoms were the fundamental particles of nature. The main ideas (postulates) were:


 * All matter consists of tiny particles called atoms.
 * Atoms can neither be created nor destroyed.
 * Atoms of an elements are indivisible and cannot be changed into atoms of any other elements.
 * All atoms of the same element are identical in properties such as size, mass etc.
 * Atoms of one element differ in properties from atoms of other elements.
 * Chemical combination is the union of atoms of different elements, forming ‘compound-atoms’ molecules).
 * In any compound the atoms are combined in constant, whole number ratios to each other.
 * Part of Dalton’s theory was that atoms of elements were solid and indivisible.

However, work carried out by many other scientists such as Faraday, Thompson, Rutherford and Bohr established that in fact, atoms consist of protons, neutron and electrons. Their discoveries led to a nuclear model of the atom and the following ideas:


 * Atoms consist of two regions, that is, a small dense nucleus surrounded by a cloud of electrons.
 * The nucleus is positively charged and contains protons and neutrons. The very large majority of the mass of an atom is contained in the nucleus.
 * The electrons are negatively charged and have a very small mass. They move rapidly in the region of space around the nucleus creating an effect of an electron cloud. This electron cloud makes up nearly all the volume of the atom.
 * Atoms are electrically neutral. Hence the number of protons is equal to the number of electrons.

The Structure of the Atom

 * The three fundamental particles which make up atoms include: protons, neutrons and electrons.




 * Protons and neutrons have approximately the same mass.
 * Electrons have a much smaller mass, about 1/2000th the mass of protons and neutrons.
 * Protons and electrons have equal but opposite charges - protons have a positive charge and electrons have a negative charge.
 * Neutrons are uncharged.

A Model of the Atom

 * From many experiments, particularly those of Lord Rutherford early this century, the atoms is believed to have a relatively small, dense nucleus consisting of protons and neutrons. The electrons are located within a region called the electron cloud which constitutes the atom’s volume.

//Diagram of an Atom//

The Structure of the Nucleus

 * The central part of the atom which contains the protons and neutrons has a positive charge equal to the number of electrons.
 * Exceedingly small when compared with an atom.
 * Contains 99.9% of the mass of an atom - due to the relatively larger masses of the protons and neutrons when compared to the masses of the electrons.
 * Dense - due to its large mass and small volume

The Arrangement of the Electrons around the Nucleus

 * Electrons move through a relatively large space outside the nucleus.
 * Kept moving around the nucleus by attractive electrostatic forces between the positively charged nucleus and the negatively charged electrons.
 * In a neutral or uncharged atom, the number of electrons equals the number of protons.

Atoms can differ in atomic number and mass number:


 * **Atomic Number** - is the number of protons in the nucleus of an atom - symbol (Z). In an electrically neutral atom, the number of protons equals the number of electrons. If an atom has no net charge (it is not an ion) the atomic number represents the number of protons and the number of electrons.


 * **Mass Number** - is the sum of the number of protons and neutrons - symbol (A). If the mass number and the atomic number for an atom is known, it is possible to work out the number of neutrons in the nucleus of that atom:


 * Mass Number(A) = Atomic Number(Z) + Number of Neutrons


 * The convention used to identify the structure of an atom (also called a **nuclide**):

//Diagram//

Isotopes

 * All atoms of a given element have the same number of protons but the number of neutrons may vary. Hence different forms of an element may have a different mass number (A). These different forms of an element are called isotopes.

//Examples//

=The Periodic Table=

//Periodic Table//


 * Elements in the period table are listed in order of increasing atomic number but in such a manner as to created vertical groups with similar chemical properties.
 * Elements are arranged in order of increasing atomic number.
 * The horizontal rows of the periodic table are called periods and the vertical lists are called groups.

Periods (horizontal rows)
the first period are H and He the second period contains eight elements - Li, Be, B, C, N, O, F, Ne
 * The horizontal rows of the periodic table are called periods. Elements contained in:

Other changes that occur include: changes in reactivity. ionisation energies etc.
 * There is a steady change in the properties of elements across the period. For example, in the 3rd period there is a decrease in the metallic properties of the elements.

Groups (vertical columns)

 * The vertical columns in the periodic table are called groups.
 * The ten columns in the middle of the periodic table are called the transition elements.
 * The groups contain elements with similar chemical properties, generally with a gradual change in physical properties:

Group 1 Elements - The Alkali Metals

 * Excluding hydrogen, contains the alkali metals. These are soft solids with low densities which increase down the group (the first three are less dense than water).
 * They are all strong reducing agents (electron donor) because of their tendency to form 1+ ions.
 * They readily tarnish in air forming an oxide crust and burn vigorously when ignited. 2Na (s) + O2 (g) → Na2O2 (s)
 * They react with halogens (Group VII) and sulphur to form ionic solids. 2Na (s) + Cl2 (g) → NaCl (s)
 * Their solution, when they react with water, are strongly alkaline. 2Na (s) + 2H2O (l) → 2( Na+ (aq) + OH- (aq) ) + H2 (g)
 * Because of their reactivity, which increases down the group, they are stored under paraffin oil or kerosene.

Group II Elements - The Alkaline Earth Metals

 * Also reactive, but less so than the alkali metals. They are also reducing agents (electron givers) because of their tendency to form 2+ ions. Thus react at higher temperatures than the alkali metals, with air, halogens, water, acids,

Group VII Elements - The Halogens

 * Are reactive non-metals. They have a tendency to form 1- ions and are thus good oxidising agents (electron acceptor). They also react with non-metals to form covalent molecular species. They also react with most metals to form ionic compounds. Mg (s) + Cl2 (g) > MgCl2 (s)


 * The reactivity of the halogens decreases down the group from fluorine to iodine. Similarly, their physical properties show trends, i.e. the melting points and boiling points increase down the group so that fluorine and chlorine are gases at normal temperatures, while bromine is a liquid and iodine is a solid.


 * Their colours also become more intense, from the pale yellow of fluorine, through the yellow-green of chlorine and red brown of bromine to the black violet of iodine.

Group VIII Elements - The Noble (Inert) Gases

 * Almost completely unreactive, noble gases only occur in the atmosphere.
 * The physical properties of the noble gases show trends - the melting points and boiling points increase down the group, with helium having a melting point of -272 C and a boiling point of -269 C. Their densities also increases in the same direction.

=Metal and Non-Metals=


 * Elements to the left and below the diagonal line are metals
 * Elements to the right and above the diagonal line non-metals.
 * Semi-metals are elements which are difficult to classify as metals or non-metals: boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb) and tellurium (Te).
 * The semi-metals located, on the period table, along the diagonal line separating the metals from the non-metals. Also referred to as metalloids.

=Bonding=

Essentially, chemical bonds are the force holding atoms, ions or molecules together. The forces are electrostatic - a balance of the attraction of oppositely charge particles and the repulsion of similarly charged particles. The types of bonds formed depend on the outer electron configuration of atoms and the nature of the atoms which are being bonded.

The structure of a substance depends on:
 * the nature of the particles present.
 * the force holding the particles together.

Elements can be classified into four groups:
 * metals
 * covalent molecular substance
 * covalent network elements
 * ionic substances

Metallic Substances
All metallic substances tend to be solid at room temperature; with exception of liquid mercury Physical characteristics that all metals share include:
 * relatively dense solids.
 * good conductors of heat and electricity.
 * malleable and ductile
 * have shiny surface when freshly cut or cleaned.
 * relatively high melting points.

The above physical characteristics have enabled scientists to make a number of inferences about the metals:
 * The high densities indicate that the atoms are packed tightly together.
 * Malleability and ductility indicates that there is only limited resistance to movement of atoms with respect to each other.
 * Electrical conductivity indicates that metal contain mobile charged particles.

Metallic Bonding
Chemists have devised a model about the structure of metals: Metals conduct electricity because of the mobility of the electrons within the lattice. Electrons entering one end of the metal cause a similar number of electrons to displaced from the other end and the metal conducts.
 * the outermost or valence electrons in metal atoms move about freely with in a three-dimensional arrangement or lattice of positively charged metal ions.
 * consists of positive ions surrounded by a ‘sea’ of electrons.
 * valance electrons are delocalised - they are not associated with a particular metal ion and move through the lattice of metal ions.
 * the negatively charged electrons are attracted to the positively charged metal ions in the lattice and this electrostatic attraction holds the lattice together - metallic bonding.

Metals contain delocalised electrons: electrons which are not associated with a particular metal ion and move through the lattice of metal ions - these mobile electrons acquire heat from a heat source and rapidly transfer it to cooler parts of the lattice. Because these electrons do not belong to any particular atom, if enough force is applied, one layer of atoms can slide over another layer without disrupting the metallic bonding.

Metals have high melting and boiling points due to strong electrostatic attraction between the positive metal ions and the delocalised electrons. Group II of the periodic table form the strongest metallic bonds because they tend to form di-positive metal ions and lose two electrons per atom.

Uses of Metals

 * Aluminium - used in domestic utensils, drinks cans, cooking foil etc. Used due to it’s thermal conductivity, malleability and attractive lustre.
 * Copper - electrical wiring; excellent conductors of electricity.

Covalent Molecular Substances
These are diatomic molecules: molecules which consist of two atoms bonded together. Other non-metal elements exist as polyatomic molecules: consist of several atoms bonded together in the molecule. Examples include - phosphorous, P4, and sulphur, S8.

In covalent molecular compounds the formula represents the number of atoms of each element in one molecule of the compound:


 * hydrogen chloride: HCl - contain 1 atoms of hydrogen and 1 atom of chlorine.
 * water: H2O - contains 2 atoms of hydrogen and 1 atom of oxygen
 * ammonia: NH3 - contains 1 atom of nitrogen and 3 atoms of hydrogen

The procedure used to name molecular compounds:
 * The name of the elements closer to the bottom or left hand-side of the periods table is written first.
 * The second part of the name is obtained by adding the suffix, ‘ide’, to the stem of the name of the second element.
 * Where a molecule contains more than one atoms of one type the number of atoms is indicated by the prefixes ‘mono’, ‘di’, ‘tri’, ‘tetra’, ‘penta’ and ‘hexa’ which stand for 1, 2, 3, 4, 5 and 6 respectively. The prefix ‘mono’ is not used for the first-named element.

Covalent Network Substances
Substance such as diamond (C) and silica (SiO2), consist of three dimensional network structures wherein all the atoms are linked together by strong covalent bonds. These covalent bonds are continuous throughout the structure and results in a very hard structure with high boiling points. Covalent network substance have the following characteristic properties:


 * Very high melting and boiling points
 * Non-conductors of electricity
 * Extremely hard and brittle
 * Chemically inert
 * Insoluble in water and most other solvents.

The properties listed above are the result of strong covalent bonding giving strength and high melting points. The highly localised electrons prevent the conduction of electricity, except for graphite which has one delocalised electron from each carbon atom.

Fullerenes - A New Form Of Carbon
A third form of carbon, made up of cage-like molecules, has only recently been discovered (1985). These molecules consist of large number of atoms covalently bonded together as a series of pentagons and hexagons and usually forming a spherical hollow shape. The most stable and interesting fullerene is C60 which is made up of 60 carbon atoms arranged as 12 pentagons, each surrounded by a hexagon. Fullerenes were named after the American architect, Buckminister Fuller, who was noted for his designs of geodesic domes of high stability. The C60 molecule was given the name buckminster fullerene but it is usually referred to as ‘buckyball’ for short. Buckyballs form a soot like substance whose properties and uses are still under research.

Ionic Substances/Compounds
The basic particles which make up ionic compounds are ions: electrically charged atoms which have lost or gained electrons.

Simple Ions

 * Cations** - positively charged atoms which are formed when one or more electrons are removed from an atom.


 * Anions** - negatively charged atoms which are formed when one or more electrons are gained by an atom.

Elements which are cations are metallic elements - exception hydrogen and ammonium. Anions are non-metal atoms.

Polyatomic Ions
Polyatomic ions are groups of atoms bonded to one another which have a net positive or negative charged.

Examples include: phosphate, PO43-, carbonate CO32-, ammonium NH4+.

The types of bonds holding the atoms within the polyatomic ion together are chemical bonds similar to those involved in molecules (i.e. covalent bonds)

The system used for naming ionic compounds:


 * The name of the positive ion is written first.
 * The name of the negative ion is written second.
 * Where the ionic compound has water molecules of crystallisation, the number of water molecules is indicated.

Ionic Substances
Ionic compounds contain oppositely charged ions which are arranged in regular three-dimensional lattice.

The physical properties of ionic compounds are:
 * Hard and brittle
 * Non-conductors of electricity in the solid state, but good conductors when molten or in aqueous solution.
 * High melting and boiling points.

The conductivity of molten ionic compounds suggests that ionic substance contain positively charged metal ions and negatively charged non-metal ions. Particles which are free to move when in the molten state.

Chemical reactions occur at the electrodes when an electrical current is passed through molten sodium chloride:
 * At the electrode connected to the positive terminal; chlorine gas forms - chloride ions move to this electrode and are negatively charged.
 * The electrode connected to the negative terminal; sodium metal is formed - sodium ions move to this electrode and are positively charged.

Electrolysis is the term used to describe the process by which direct current is passed through molten substances or conducting solutions to produce chemical reaction at the electrodes.

Ionic Bonding
The structure of an ionic solid can be best described by using an example such as sodium chloride:


 * ions in the crystal are arranged in a regular three-dimensional lattice.
 * each sodium ion is surrounded by six negative chloride ions and each negative chloride ion is surrounded by six positive sodium ions.
 * when a solid, the position of the ions is fixed.

In ionic solids, are ions held in the crystal lattice by strong electrostatic attraction to the oppositely charged ion around it - called ionic bonding.

Strong attractive force exist between oppositely charged ions; if a layer of crystal is forced to slide past another layer, the orderly arrangement of ions is disturbed. Ions of similar charge are forced closer to another with an increase in repulsive forces and decrease in attractive forces - the crystal fractures.

This makes ionic compounds:
 * hard
 * brittle, and
 * difficult to cut

Ionic solids have high melting and boiling points because the forces of attraction between oppositely charged ions in ionic compounds are very strong.

An ionic solid does not conduct electricity; the ions are in fixed positions and can not move. However, molten ionic compounds and ionic compounds dissolved in water conduct electricity because the ions are free to move. They do not conduct electricity as well as metals.

Ionic compounds form between metal and non-metal elements. All substances formed when metals from groups I and II of the periodic table react with non-metals from groups VI and VII are ionic compounds.

=Properties, Structure and Bonding of Elements and Compounds= = =

=Electron Configuration of Ions=

Electron Arrangement in Atoms
The way electrons are arranged in atoms is very important as it determines chemical behaviour. In 1912 the Danish Physicist Niels Bohr proposed a theory of the atom which was able to explain more clearly the behaviour of electrons in atoms and was consistent with the quantum theory. He proposed that:
 * electrons can only exist in specific energy levels.
 * electrons could be excited from one level to another by specific amounts of energy corresponding to the difference in energy levels.
 * energy, in form of photons, is emitted whenever an electron moves from a high energy level to another.

The Bohr model proposed a central, dense positively charged nucleus, with electrons orbiting the nucleus in fixed, circular trajectories, rather like the solar system of sun and planets.

The permissible orbits provide discrete energy levels, or shells for electrons around the nucleus. These shells were given numbers, principal quantum numbers (n), 1, 2, 3, 4 ..... or letters K, L, M, N .... In other words, electrons existed or could be found at distinct energy levels or shells around the nucleus of atoms.

Shells are further divided into subshells. These subshells offer different regions of space and have different shapes. They are designated s, p, d, f ..... The subshells are further split into orbitals which may have different energy levels.

Orbitals exist within these subshells: areas within a subshell in which there is some probability of finding an electrons. Electrons do not have a definite path or orbit. The number of each type of orbital is given by 1, 3, 5, 7 .....

For a particular hydrogen atom, its one electron can occupy any of these orbitals at a particular instant. When it is in the lowest possible energy state (occupying the 1s orbital) the atom is said to be in the ground state. If the electron absorbs enough energy to move to a higher energy orbital, the atom is said to be in an excited state.

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Electron Configuration of Multi-Electron Atoms
This normally applies to atom in their ground state (i.e. lowest energy state).

The energy level diagrams for elements of atomic numbers greater than 2 are similar to that for hydrogen, but there is an upward displacement of the subshells, due to electron repulsions. Thus the order of increasing energy show significant changes to that of hydrogen.

Two electrons occupy each orbital.

The sequence of orbital filling or increasing energy is as follows:

So the electron filling sequence for potassium would be:

19K (19 electrons so) 1s2 2s2 2p6 3s23p6 4s1 (2, 8, 8, 1)

Why elements bond
The Octet Rule - having a valence electron configuration of eight electrons - all noble gases, other than helium, have an outer energy level or valence electron configuration of eight electrons - an octet configuration. Elements attain an octet configuration when they bond.

An ionic bond is formed when elements that bond together attain a noble gas electron configuration by gaining or losing electrons.

Example: NaCl

Na electron configuration: 1s2 2s2 2p6 3s1 (2, 8, 1) is one atom more than the neon configuration of 1s2 2s2 2p6 (2, 8). So sodium loses one electron and attains the stable electron configuration of neon.

Cl electron configuration: 1s2 2s2 2p6 3s2 3p5 (2, 8, 7) is one electrons short of the stable argon electron configuration of 1s2 2s2 2p6 3s2 3p6 (2, 8, 8). When chlorine reacts with sodium it gains one electron from sodium to attain the stable electron configuration of argon.

For electrons to be removed from an atom, energy is needed due to electrostatic attractive forces. This energy is called the ionisation energy. The higher the ionisation energy, the less likely an atom is to give up electrons. Metals tend to lose electrons because they have low ionisation energies, while non-metals tend to gain electrons since they have high ionisation energies.

=Electron Dot Diagrams=

Electron dot diagrams are a simple way of showing the electron arrangements of atoms and the changes that occur as bonds are formed.

Electron dot diagrams use a dot for each valance electron. The dots are spread around the atom and, where possible, are in pairs. Consider the element aluminium and nitrogen. Aluminium has three valance electrons and nitrogen has five.

=Covalent Molecular Substances=

The bonding that exists within covalent molecular substances (e.g. iodine):
 * exist as groups of neutral atoms called molecules.
 * force between molecules are weak; indicated by low melting points.
 * large amount of energy is required to separate iodine molecules into individual atoms:

I2 (g) → 2I (g)

Strong bonding forces exist between iodine atoms within the iodine molecules.

Covalent molecular substance do not conduct electricity due to the lack of charged species in the solid and liquid states - there are no delocalised electrons nor charged ions capable of conducting an electric charge.

Covalent Bonding
Covalent bond - occurs when atoms in a molecule share electrons and are held together by the electrostatic attraction between the shared electrons and the nuclei of adjacent atoms.

Example: Hydrogen gas

H atom + H atom → H2 molecule


 * both hydrogen atoms have an electron configuration of 1s1.
 * both need one electron to attain the electron configuration of helium.
 * both have the same ionisation energies so electron transfer can not occur.
 * instead each H atom shares it single 1s1 electron - each atom has the stable electron configuration of He.

Neon and other noble gases exist as simple atoms since all the orbitals of neon are fully occupied, so no covalent bond formation is possible.

Covalent Bonds and Valence Electrons
The number of covalent bonds formed by an element depends on the number of electrons in its valence or outer energy level.


 * Fluorine will form one single covalent bond - needs one electrons to have a share of 8 valence electrons.
 * Oxygen will form two covalent bonds - needs two electrons to have a share of 8 valence electrons.
 * Nitrogen will form three covalent bonds - needs three electrons to have a share of 8 valence electrons.
 * Carbon will form four covalent bonds - needs four electrons to have a share of 8 valence electrons.

The electron pairs forming the covalent bond are called bonding electron pairs.

The remaining electron pairs which are not involved in the formation of covalent bonds are called non-bonding electrons pairs or lone pairs.

The number of bonding electrons pair and non-bonding electron pairs are shown below in the water molecule:

Multiple Covalent Bonds
A double covalent bond - a covalent bond in which there is two shared pairs of electrons:


 * each oxygen atom is two electrons short of a valence electron octet.
 * each oxygen atom accepts a share in two electrons belong to the other oxygen atom.
 * 4 electrons are shared by the two oxygen atoms and occupy the valence energy level of both.

A double covalent bond can be represented by 2 parallel lines joining the atoms:

The type of covalent bond that forms between nitrogen atoms is a triple covalent bond, since with nitrogen there is three valance electrons short of a valance electron octet.

NOTE: In dot diagrams you can also use a ‘x’ to represent the electrons from other attached elements.

Coordinate Covalent Bonds
When elements in groups V, VI and VII form molecules, they often have non-bonding pairs of electrons. These atoms can form additional covalent bonds by sharing their lone electron pair with other atoms which have completely vacant valence orbitals. Since both electrons in the shared pair originate from one atom the bond formed is called a coordinate covalent bond.

The method used to represent coordinate covalent bond is an arrow since it distinguishes the coordinate bond from other covalent bonds; both type bonds are identical with respect to properties:

Example: Ozone, O3: coordinate covalent bond

Some Non-Octet Molecules
The octet rule is often broken by atoms whose valance energy levels have a higher electron capacity than eight electrons. This is the case in atoms of elements in period 3 and beyond.

For example, phosphorous has 5 valence electrons in its 3rd energy level - so when it form PCl5, it share one electron from each of the five chlorine atom which gives a total of five bonding pairs in the outer energy level:

Compounds With Both Ionic And Covalent Bonding
Some compounds contain both ionic and covalent bonds. For example, when ammonia gas is mixed with hydrogen chloride gas, white clouds of ammonium chloride form:

NH3 (g) + HCl(g) → NH4Cl (s)

The bonds between the ammonium ion and the chloride ion is ionic, but the bonds within the ammonium ion is covalent.

=An Introduction to Hydrogen Bonding=


 * Commonly occurs in NH3, H2O and HF
 * Polar molecule (one side more positive than other)
 * Forms stronger intermolecular bonds than expected due to dipole nature of molecule (two distinct polar ends – positive and negative)
 * Line up so that positive and negative ends of different molecules are together – dipole attraction


 * This dipole attraction is called “hydrogen bonding” and is shown in water


 * Hydrogen bonding accounts for many special properties of water e.g. increase in volume when water freezes.