Atomic Bonding

1.1
The atoms, during bond formation, may lose or gain electrons (valence electrons) in order to achieve a stable state, or technically speaking, a stable electron configuration. Usually metal atoms lose electrons and non-metals gain electrons in order to achieve electron stability. When dealing with bond formation (Ionic bonding for example) we need to analyse the outer shell of the atom. Metals usually present 1, 2 or 3 electrons in their outer shell therefore they have to give them away to achieve stability. Vice versa non-metals have 5, 6 or 7 electrons in their outer shell and they need to receive more electrons in the outer shell to be stable.
1.2
Ionic bonding occurs between metal and non-metal atoms and consists in gaining and losing electrons to reach electron stability. Metal atoms lose electrons (negative charge) becoming ions, positively charged. Non-metal atoms gain electrons to achieve electron stability.
Sodium has one valence electron in its outer shell, so it needs to give that electron away in order to achieve stability. Vice versa, Chlorine atom has 7 valence electrons in its outer shell and it needs to gain one electron (from sodium atom) to achieve electron stability.

1.3
Covalent bond occurs between non-metals and consists in a shared pair of electrons, from the outer shell, between two atoms. Each of the two atoms, involved in the bonding, provides an electron. Usually non-metals present 4 or more electrons in their outer shell and a lot of energy is needed to remove electrons in order to form bonds. That’s one of the reasons why covalent bonding is used between non-metals. The electrons shared are called molecules.
When the pair of electrons is shared from only one of the two atoms involved in the reaction, then it’s called dative covalent bond. It is represented by a short arrow going from the electron giving both electrons towards the one providing neither.

2.1

Metallic bond is a strong electrostatic attraction between a sea of delocalised electrons and a lattice of positive ions, called cations. This strong electrostatic attraction holds the structure together, avoiding that positively charged ions repel the negatively charged electrons.
Because of the fact that electrons are free to move metals have the characteristic to be good conductor of heat and electricity. The metals are also ductile, which it means that they can be stretched out to make a wire. They are malleable too. This means that they can pressed and beaten to change their shape.

http://www.ndt-ed.org/EducationResources/CommunityCollege/Materials/Structure/metallic.htm
3.1
Before defining the van der Waals’ forces we need to talk about intermolecular forces. These are forces holding molecules together and to melt or boil a substance we have to overcome these forces. The strongest these forces are, and the more energy will be required to break the substance’s bonds in order to reach its melting or boiling point.
The weakest of the intermolecular forces are called van der Waals’ forces, which are forces of attraction between a temporary dipole on one molecule, and a induced dipole on an adjacent molecule.
Temporary diploes are a surplus of electrons in one of the atoms, making it to appear asymmetrical in only a very short time because the electrons are in constant motion. If we consider an adjacent molecule then we can notice that the electrons are repelled by the negative side of the dipole and attracted by the positive side.

“Mill Hill County High School”, Scholarly articles on Atomic Bonding found on Google.

3.2

Hydrogen Bonding is the strongest intermolecular force and it can be defined as an attractive force between the Hydrogen (bonded to oxygen, nitrogen, or fluorine) and an electronegative atom of an adjacent molecule.
This happens because The Hydrogen has almost null electron density and this make it able to bond with electronegative atoms on neighbouring molecules, forming a strong intermolecular dipole-dipole bond.
Intermolecular hydrogen bond is different from the intramolecular one when the attraction between hydrogen and a electronegative atom is within the molecule.

Hydrogen bonding gives to water the characteristics of a high boiling point, low vapour pressure and a low density in its solid form, occupying more volume. The Hydrogen bonding in water is formed between the electronegative oxygen atom and Hydrogen atoms. The electronegativity of the oxygen atoms is due to the fact that it is surrounded by two bonding pairs and two lone pairs. This unshared pairs gives to the oxygen the partial negative charge.

http://www.biology.arizona.edu/biochemistry/tutorials/chemistry/page3.html

The hydrogen bond is really a special case of dipole forces. A hydrogen bond is the attractive force between the hydrogen attached to an electronegative atom of one molecule and an electronegative atom of a different molecule. Usually the electronegative atom is oxygen, nitrogen, or fluorine, which has a partial negative charge. The hydrogen then has the partial positive charge.

Ammonia with hydrogen bonding:

To recognize the possibility of hydrogen bonding, examine the Lewis structure of the molecule. The electronegative atom must have one or more unshared electron pairs as in the case of oxygen and nitrogen, and has a negative partial charge. The hydrogen, which has a partial positive charge tries to find another atom of oxygen or nitrogen with excess electrons to share and is attracted to the partial negative charge. This forms the basis for the hydrogen bond.
Ammonia molecule has nitrogen in the centre surrounded by three hydrogen atoms, l. Think (or draw) in three dimensions. The atom of nitrogen is more electronegative than the hydrogen. When the hydrogen bonding occurs, a hydrogen (with positive partial charge) of one molecule of ammonia will be attracted to the nitrogen atom, with electrons to share of the adjacent molecule.

4.1
The matter can exist in one of the three state; solid, liquid and gas. Each of the state has different physical propriety as shape, volume, and etc.

Solid In a solid the particles are tightly close and held together in lattice structure. A lattice is a systematic, regular repeating of the particles’ arrangement.
Solid can only vibrate about a fixed point and they, also, have a fixed shape and volume. The particles, although packed closely together, are in constant kinetic motion.
If enough heat energy is applied to the solid, the movement of the particles increases, starting to move and vibrate faster. The bonds, between the particles, end up being weaker and at the melting point the heat energy is used to break some of the bonds instead of increasing the solid’s temperature. At this point a liquid is formed.

https://thescienceclassroom.wikispaces.com/Taylor's+States+of+Matter

Liquid
The particles in a liquid are still close together, although there is enough space between them to make them moving faster. At this stage the liquid has a fixed volume but not shape, and can rotate and vibrate. The liquid takes up its container shape.
When enough heat energy is supplied to the liquid the particles start to move faster , leading the weak bonds to be broken. At this point the particles are free to move far apart from each other and, finally, a gas is formed.

http://chemstuff.co.uk/academic-work/year-7/particle-model-of-solids-liquids-and-gases/image138/

Gas
The particles are free to move far apart randomly. The kinetic energy of the particles is bigger than the force of attraction between them. A gas has neither a fixed shape nor a fixed volume.

http://www.mit.edu/~kardar/teaching/projects/chemotaxis(AndreaSchmidt)/more_random.htm

4.2

If we have a gas inside a container with a fixed volume, then we can notice that if we apply heat energy to the container the gas particles will gain kinetic energy, moving faster. This speed will make them colliding more frequently with the container walls. This will cause that the forces acting on the container walls will increase and so the pressure. This is known as gas pressure and states:

The pressure is proportional to the temperature if the gas temperature is measured in Kelvin scale.

P/T = V constant P = pressure
P1/T1=P2/T2 V = volume T = temperature
4.3

5.1

Valence Shell Electron Pair Repulsion Theory (VSEPR) predicts the geometry or shape of each atom in a molecule, minimising the electrostatic forces of repulsion between electrons in the valence shell of that atom. These are the five main molecular shapes:

Water (H2O) has a bond angle of 104.5° because the two lone pairs around the oxygen repel the two bonding pairs so strongly. This leads the bond angle of water to be even smaller than a tetrahedral electron pair geometry, which has a bond angle of 109°. The water has a tetrahedral shape as arrangement of the pairs of electrons and a bent line as shape of the molecule.

http://www.elmhurst.edu/~chm/vchembook/206bent.html

Ammonia (NH3) has three bonding pairs and one lone pair around the nitrogen atom. The distance between the four pairs is 109° (Tetrahedral electron pair geometry). The lone pair repels the bonding pairs more strongly that the bonding pairs do between each other. This repulsion reduces the bond angle to 107°.

http://www.elmhurst.edu/~chm/vchembook/205trigpyramid.html

Carbon Dioxide (CO2) has two double bonding pairs around the central carbon, which act as single bonds. The two double bonding pairs repel each other so far away that the molecule shape results to be linear with a bond angle of 180°.

http://www.elmhurst.edu/~chm/vchembook/202linear.html

Methane (CH4) has a tetrahedral shape because it has a central carbon atom, surrounded by four bonding pair of Hydrogen electrons. The four pairs are all bonding pairs and therefore they don’t exert a high repulsion force. The bong angles are all 109.5°.

http://www.elmhurst.edu/~chm/vchembook/204tetrahedral.html
Boron Trifluoride (BF3) has central Boron atom surrounded by three bonding pairs of fluorine electrons, forming a trigonal planar shape. The three bonding pairs exert a minimum repulsion force and the bond angle between each bonding pair is 120°.

The sulphur Hexafluoride has an octahedral shape, in which the sulphur atom is surrounded by six bonding pairs of fluoride electrons. This has the shape of two Pyramid attached bas to base.The distance between each bonding pair is 90°.

http://en.wikipedia.org/wiki/File:Sulfur-hexafluoride-2D-dimensions.png http://www.funcrunch.com/sulfur-hexafluoride-8385/

5.2

The sodium chloride (NaCl) structure is an ionic structrure, where a large number of sodium and chloride ions arranged in regular lattice, forming a crystal.

The stable form of sodium chloride involves a very large number of NaCl units arranged in a lattice (or regular arrangement) millions of atoms across. Because the lattice is rigid, this means that one gets a solid: the ions do not move much one with respect to another. Also, because atoms are so small, even a small crystal of salt will have billions of sodium chloride units in it! The ions are arranged so that each positive (sodium) ion is close to many negative (chloride) ions, as shown on the following picture

The ions in a compound such as sodium chloride are arranged in a lattice structure. This regular arrangement results in the formation of a crystal. This arrangement shows that each ion is surrounded by ions of opposite charge. Sodium chloride crystal has giant three-dimensional lattice structure, as shown in the diagram below, because there are tiny sodium chloride structure repeated billions or even trillions of times. There are strong electrostatic forces between the positive and negative ions, giving to the sodium chloride the shape of a strong structure with a high melting and boiling point. This structure gives to the ionic compound the ability to not conduct electricity in the solid state. It can conduct electricity only if melted of dissolved in water.

http://www.bbc.co.uk/schools/gcsebitesize/science/add_edexcel/ionic_compounds/ionicrev5.shtm http://www.chm.bris.ac.uk/pt/harvey/gcse/ionic.html Diamond is an allotrope of carbon, having the same state but different structure. Diamond has a giant covalent lattice structure in which each carbon atom is joined to four other carbon atoms in a tetrahedral arrangement, forming four covalent bonds. It is a giant covalent structure because the arrangement mentioned above is repeated very many times. The diamond structure is very strong, therefore it has a high melting point and does not conduct electricity.

http://www.uwgb.edu/dutchs/Petrology/Diamond%20Structure.HTM

Element | Symbol | Electronic configuration | Hydrogen | H | 1s1 | | Helium | He | 1s2 | Lithium | Li | [He] 2s1 | Beryllium | Be | [He] 2s2 | Boron | B | [He] 2s2 2p1 | Carbon | C | [He] 2s2 2p2 | Nitrogen | N | [He] 2s2 2p3 | Oxygen | O | [He] 2s2 2p4 | Fluorine | F | [He] 2s2 2p5 | Neon | Ne | [He] 2s2 2p6 | Sodium | Na | [Ne] 3s1 | Magnesium | Mg | [Ne] 3s2 | Aluminium | Al | [Ne] 3s2 3p1 | Silicon | Si | [Ne] 3s2 3p2 | Phosphorus | P | [Ne] 3s2 3p3 | Sulfur | S | [Ne]3s2 3p4 | Chlorine | Cl | [Ne] 3s2 3p5 | Argon | Ar | [Ne] 3s2 3p6 | Potassium | K | [Ar] 4s1 | Calcium | Ca | [Ar] 4s2 | Scandium | Sc | [Ar] 3d1 4s2 | Titanium | Ti | [Ar] 3d2 4s2 | Vanadium | V | [Ar] 3d3 4s2 | Chromium | Cr | [Ar] 3d5 4s1 | Manganese | Mn | [Ar] 3d5 4s2 | Iron | Fe | [Ar] 3d6 4s2 | Cobalt | Co | [Ar] 3d7 4s2 | Nickel | Ni | [Ar] 3d8 4s2 | Copper | Cu | [Ar] 3d10 4s1 | Zinc | Zn | [Ar] 3d10 4s2 | Gallium | Ga | [Ar] 3d10 4s2 4p1 | Germanium | Ge | [Ar] 3d10 4s2 4p2 | Arsenic | As | [Ar] 3d10 4s2 4p3 | Selenium | Se | [Ar] 3d10 4s2 4p4 | Bromine | Br | [Ar] 3d10 4s2 4p5 | Krypton | Kr | [Ar] 3d10 4s2 4p6 |
5.4

5.5 Ionisation energy is the energy necessary to pull away an electron in the outer shell of a gaseous atom, leaving a positively charged ion. The ionisation is present only in gas because the atom is by itself and not attached to others as it happens in solids and liquids. The term ionisation comes from the ion that it is created from the process.

The factors that influence the ionisation energy are: * Nuclear charge

If the nuclear charge increases the ionisation energy increases as well because the attraction force between the nucleus and the valence electrons becomes stronger.

* Size of the atom

The Ionisation energy is influenced by the size of the atom. If the atom increases the attraction force between the nucleus and the valence shell decreases and vice versa.

* Penetration effect

The Ionisation energy is stronger near the nucleus, so in the sub level s is stronger than in the p, in the p is stronger than in the d and so on.

* Electronic configuration

The ionisation energy is influenced by the electronic configuration because it is higher in atoms with a stable electronic configuration, where there is less probability to electrons.