Atomic+Structure+and+Bonding


 * __Atomic Structure and Bonding__ **

**- Atoms of different elements differ in size and mass **
 * __Atoms__ **
 * - Smallest particle of any substance **

__**3 Types of Sub-atomic particles: ** __**2. Neutrons (In nucleus) **__ __**3. Electrons (Around nucleus) **__
 * 1. Protons (In nucleus) **__
 * Charge : +1 **
 * Mass : 1 a.m.u **
 * Charge : 0 **
 * Mass : 1 a.m.u **
 * Charge : -1 **
 * Mass : 0.00055 (1/1800) a.m.u **
 * (a.m.u = atomic mass unit) **


 * <span style="font-family: 'Arial','sans-serif'; font-weight: normal;">- Outer part of the atom that is involved in chemical reactions. Therefore, electrons are the most important particles because they surrounds the nucleus and therefore are in the outer part of the atom. **

__**<span style="font-family: 'Arial','sans-serif';">Atomic Number (Z): **__
 * <span style="font-family: 'Arial','sans-serif'; font-weight: normal;">- Number of protons in the atom **
 * <span style="font-family: 'Arial','sans-serif'; font-weight: normal;">- In a neutral atom, number of protons = number of electrons, therefore, Atomic Number = No. of Protons = No. of Electrons **
 * <span style="font-family: 'Arial','sans-serif'; font-weight: normal;">- Atomic number is the same for every atom of the element. Therefore, the chemical identity of the atom can be determined from the atomic number **
 * <span style="font-family: 'Arial','sans-serif'; font-weight: normal;">- Mass Number = No. of protons + No. of neutrons **

__**<span style="font-family: 'Arial','sans-serif';">Proton and Mass Numbers: **__
 * <span style="font-family: 'Arial','sans-serif'; font-weight: normal;">23 (Mass Number) **
 * <span style="font-family: 'Arial','sans-serif';">Na **
 * <span style="font-family: 'Arial','sans-serif'; font-weight: normal;">11 (Proton Number) **


 * <span style="font-family: 'Arial','sans-serif'; font-weight: normal;">- Proton Number = No. of protons in the atom **
 * <span style="font-family: 'Arial','sans-serif'; font-weight: normal;">Example: Sodium has a proton number of 11. Hence, every sodium atom will have 11 protons and every atom with a proton number of 11 is a sodium atom. **


 * <span style="font-family: 'Arial','sans-serif'; font-weight: normal;">- Mass Number = No. protons + No. of neutrons **
 * <span style="font-family: 'Arial','sans-serif'; font-weight: normal;">Hence, number of neutrons = mass number - proton number **

NOTE: Atomic number is the smaller of the two numbers in the Periodic table Isotopes <span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%;">- Atoms that have the same atomic number but different mass numbers - Example: 1H, 2H, 3H - Relative atomic mass of element is the average masses of the isotopes. Example: 25% is 35 while 75% is 37. So, (25% x 32 + 75% x 37) / 100= atomic mass of element - Similar to the element itself because: ~ Isotopes have similar chemical properties because they have the same number of electrons [because they have the same electron configurations.] (Electrons are involved in the chemical processes) ~ But, they have different physical properties because they have a different mass. e.g Mass no. 12 Atomic no. 6 || <span style="color: #000000; display: block; font-family: Arial,sans-serif; font-size: 11.5pt; text-align: center;">6 || <span style="color: #000000; display: block; font-family: Arial,sans-serif; font-size: 11.5pt; text-align: center;">6 || <span style="color: #000000; display: block; font-family: Arial,sans-serif; font-size: 11.5pt; text-align: center;">6 || Mass no. 13 Atomic no. 6 || <span style="color: #000000; display: block; font-family: Arial,sans-serif; font-size: 11.5pt; text-align: center;">6 || <span style="color: #000000; display: block; font-family: Arial,sans-serif; font-size: 11.5pt; text-align: center;">7 || <span style="color: #000000; display: block; font-family: Arial,sans-serif; font-size: 11.5pt; text-align: center;">6 || Mass no. 14 Atomic no. 6 || <span style="color: #000000; display: block; font-family: Arial,sans-serif; font-size: 11.5pt; text-align: center;">6 || <span style="color: #000000; display: block; font-family: Arial,sans-serif; font-size: 11.5pt; text-align: center;">8 || <span style="color: #000000; display: block; font-family: Arial,sans-serif; font-size: 11.5pt; text-align: center;">6 ||
 * <span style="color: #000000; display: block; font-family: Arial,sans-serif; font-size: 11.5pt; text-align: center;">Isotope || <span style="color: #000000; display: block; font-family: Arial,sans-serif; font-size: 11.5pt; text-align: center;">No. of protons || <span style="color: #000000; display: block; font-family: Arial,sans-serif; font-size: 11.5pt; text-align: center;">No. of neutrons || <span style="color: #000000; display: block; font-family: Arial,sans-serif; font-size: 11.5pt; text-align: center;">No. of electrons ||
 * <span style="color: #000000; display: block; font-family: Arial,sans-serif; font-size: 11.5pt; text-align: center;">Carbon
 * <span style="color: #000000; display: block; font-family: Arial,sans-serif; font-size: 11.5pt; text-align: center;">Carbon
 * <span style="color: #000000; display: block; font-family: Arial,sans-serif; font-size: 11.5pt; text-align: center;">Carbon

<span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%;">__**Arrangement of electrons in the atom**__ - Electrons move in electron shells. - Each shell has a limit to the number of electrons it can hold. First shell: 2 Second shell: 8 Third shell: 8 (continues) - Maximum number of electrons in the third shell is 8 for the first twenty elements in the periodic table. - Main Rule: Electrons always go to the shell nearest to the nucleus. Only when the shell has full valence will the electrons then fill the space in the next shell. - No. of electrons = No. of protons (Can be seen from the diagram provided)

Valence shell and Valence electrons - Valence shell = Outer shell = Shell furthest away from the nucleus - Valence electrons = Electrons in Valence shell - Chemical properties depend on the number of electrons present in the valence shell. Therefore, there is a correlation between the electronic configuration and the chemical behavior

<span style="color: #96e250; font-family: Arial,Helvetica,sans-serif; font-size: 132%;">**__Element__** <span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%;"> - A basic substance that cannot be simplified

<span style="color: #96e250; font-family: Arial,Helvetica,sans-serif; font-size: 132%;">**__Molecules__** <span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%;"> - Formed when two or more atoms join together chemically - Not all molecules are compounds - Example: Hydrogen (H2), Oxygen (O2), Nitrogen (N2)

<span style="color: #96e250; font-family: Arial,Helvetica,sans-serif; font-size: 132%;">**__Compound__** <span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%;"> - Contains molecules that contain at least two different elements - All compounds are molecules - Example: Water (H2O), Carbon Dioxide (CO2), Methane (CH4)

<span style="color: #96e250; font-family: Arial,Helvetica,sans-serif; font-size: 132%;">**__Periodic Table__** <span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%;"> - Position of an element is determined by its proton number - Vertical rows = Groups = Elements have the same number of valence electrons - Horizontal rows = Period = Elements have the same number of electron shells

<span style="color: #96e250; font-family: Arial,Helvetica,sans-serif; font-size: 132%;">**__Ions__** <span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%;"> - Ions are atoms with a charge, either positively or negatively charged - Number of protons =/= Number of Electrons

- Most non-metals have more than 3 electrons in their outer shell - Usually gains electrons to form negatively-charged ions - Number of electrons > Number of protons
 * __Anions__**

- Metals in group 1 and 2 usually have only 1 or 2 electrons in their outer shell - Usually lose electrons to form positively-charged ions - Number of protons > Number of electrons
 * __Cations__**

**__Types of chemical bonding__** <span style="font-family: 'Arial','sans-serif'; font-size: 11.5pt; line-height: 115%;"> <span style="font-family: Arial,sans-serif;">__**Ionic Bonding**__ - Metals and Non-metals - Electrons are transferred from metals to non-metals: Metal gives, forms a positive ion (cations) Non-metal gains, forms negative ion (anions) - Both atoms then achieves the stable electronic configuration - In the compound, ions attract ions of opposite charged (Negative attracts Positive and vice versa) - The cations and anions will then be held together by the electrostatic attraction/force in an ionic compound because of the opposite charged ions - The ions will form a giant lattice - Electronegativity : Measure of an atom's attraction for electrons ~ Electronegative = Atoms having greater tendency to attract electrons ~ Electropositive = Atoms having greater tendency to lose electrons

Can be represented using Dot-and-Cross diagrams: ~ Dots and crosses in the diagrams represent only the valence electrons (electrons from the outermost shell)

<span style="font-family: Arial,sans-serif;">__**Covalent Bonding**__ - Non-metals and non-metals - Electrons are shared between the atoms of the elements The pair of shared electrons are called covalent bond - Electrons bond to form covalent compounds (Compounds that only contain covalent bond) - Involves only the valence electrons - Each electron in a shared pair is attracted to the nuclei of both atoms, hence this is the force of attraction that hold the two atoms together. - Forms a neutral molecule (All the atoms have complete valence shells) with a noble gas configuration - Atoms can form more than a single bond to achieve noble gas configuration: - Two pairs of electrons in a covalent bond = Double bond (two lines to represent bond in diagram) - Three pairs of electrons in a covalent bond = Triple bond (three lines to represent bond in diagram)

Can be represented by Dots-and-Cross diagrams

- Metals and Metals - Electrostatic force between the negatively charged delocalized electrons and the positive metal ions - Metals cannot achieve the noble gas configuration by sharing electrons between each other (covalent bond) or by transferring electrons from one to another (ionic bond) - Metals lose their valence electrons and from a lattice of regularly spaced positive ions - The valence electron lost from each atom will become a pool of electrons that move randomly throughout the lattice in the spaces between the positively charged metal ion - Valence electrons lost = Does not belong to any metal ion = Delocalized electrons
 * __Metallic Bonding__**

**__Ionization Energy__** <span style="font-family: 'Arial','sans-serif'; font-size: 11.5pt; line-height: 115%;"> - Amount of energy required to move an electron from an isolated atom in a gaseous phase - Magnitude of Ionization energy = Measure of how tightly the electron is held in the atom - The higher the ionization energy, the more tightly the electron is held in the atom, the more difficult it is to move the electron - Electrons are removed from the outermost shell to the inner shells

First Ionization Energy - The energy required to remove one mole of electrons from each mole of atoms in the gas phase to form a mole of singly charge ions in the gas phase - Equation : M(g) -> M+ (g) + e

__Second Ionization Energy__ - The energy required to remove one mole of electrons from each of a mole of singly positively charged ions in the gas phase to form a mole of doubly positively charged ions in the gas phase. - Equation : M+ (g) -> M^2+ (g) + e

- Second ionization energy is greater than the first ionization energy. This is because electrons are now being removed from positively charged cations which are able to hold onto the negatively charged electrons more strongly. - The higher the ionization energy, the more difficult it is for the electrons to be removed and hence they must be nearer to the nucleus - Ionization energy increases as successive electrons are removed. This is because once an electron has been removed from an atom/ion, the remaining electrons will be more strongly attracted to the nucleus - There will be a sharp increase in the graph between the 1st and 2nd ionization energy. This is because the 2nd ionization energy involves removing an electron from a filled inner shell, hence because the shell is closer to the nucleus, more energy is required to remove the electron. __Atomic Structure__

**__Properties__** <span style="font-family: Arial,Helvetica,sans-serif; font-size: 120%;">** __Molecular Structures__ ** <span style="font-family: Arial,Helvetica,sans-serif; font-size: 108%;"> - Contain individual discrete units

<span style="font-family: Arial,Helvetica,sans-serif; font-size: 120%;">** __Simple Molecular__ ** <span style="font-family: Arial,Helvetica,sans-serif; font-size: 108%;"> - Some non-metal elements and some compounds between non-metals - Simple atoms or molecules - Example: CO2, H2O, H2, CH4

__Types of forces present__ - Weak intermolecular forces between molecules but strong covalent bonds between atoms within each molecule __Volatility__ - Low Melting Point - Low Boiling Point Reason: The attraction between the molecules are weak intermolecular forces, thus the heat energy required to overcome them is low. [Note: Only weak intermolecular forces are overcome during the phase change. The strong covalent bonds between the atoms in a molecule are not readily broken down by heat, thus the molecules remain intact when the substance changes state.] __Hardness__ - Soft __Electrical Conductivity__ - Do no conduct electricity Reason: Made up of neutral molecules, thus there are no mobile charged particles to carry charges and conduct electricity. Exceptions: Covalent molecules like HCI, HNO3, H2SO4 can dissolve in water to form acidic solutions containing free ions, thus allowing an electric current to pass through __Solubility in water__ - Most covalent substances do not dissolve in water - But they dissolve in organic non-polar solvents

<span style="font-family: Arial,Helvetica,sans-serif; font-size: 120%;">** __Macromolecular__ ** <span style="font-family: Arial,Helvetica,sans-serif; font-size: 108%;"> - Large molecules - Long-chains - Polymers - Examples: polythene, nylon

__Types of forces present__ - Strong covalent bonds between the atoms within each molecule - Relatively weak intermolecular forces between molecules __Volatility__ - Moderate melting point - Moderate boiling point Reason: The polymer molecules are held by intermolecular forces. The larger polymer molecules allow for more points of contact with neighboring molecules, leading to more intermolecular forces between the molecules. Hence their melting point is higher as compared to substances with simple molecular structures. __Hardness__ - Variable __Electrical conductivity__ - Do not normally conduct electricity Reason: Made upon neutral molecules, thus there are no mobile-charged particles __Solubility in water__ - Usually insoluble in water - Some may be soluble in organic solvents (E.g Nailpolish remover)

<span style="font-family: Arial,Helvetica,sans-serif; font-size: 17px;">__** Giant Lattice **__

<span style="font-family: Arial,Helvetica,sans-serif; font-size: 120%;">** __Ionic Lattice__ ** <span style="font-family: Arial,Helvetica,sans-serif; font-size: 108%;"> - Ionic compounds - Examples: NaCl, MgO, KF, CaSO4

__Types of forces present__ - Strong ionic bonds - electrostatic forces between oppositely charged ions __Volatility__ - High melting point - High boiling point Reason: Ions are held in the lattice by strong electrostatic forces of attraction between the positive and negative ions 9ionic bonds). These ionic bonds are very strong and a lot of energy are required to overcome these forces. Application: Ionic compounds are thus used as refractory materials (heat resistant materials with high boiling points E.g MgO - line inside of furnaces, Al2O3 - used inside spark plugs __Hardness (Compare with metallic lattice)__ - Hard - Brittle (when force is applied, they tend to shatter into many smaller fragments) __Electrical conductivity__ - Conduct electricity only when in molten or aqueous form Reason: In molten or aqueous (dissolved in water) states, ions can move freely and hence act as charge carriers to conduct electricity. -Non-conductors in solid state Reason: In solid state, the ions are held in fixed positions in the lattice. The ions are not mobile, and thus cannot conduct electricity. __Solubility in solvents__ - Most ionic compounds can be dissolved in water, a polar solvent but insoluble in non-polar organic solvents Why are ionic compounds soluble in water but not organic solvents? Water molecules are polar molecules and thus can interact with the ions and weaken the ionic bonds. This causes the ions to separate, hence compound dissolves. Non-polar organic solvents do not interact with ions and hence are unable to weaken the strong electrostatic forces of attraction between the oppositely charged ions.

<span style="font-family: Arial,Helvetica,sans-serif; font-size: 120%;">** __Covalent Network__ ** <span style="font-family: Arial,Helvetica,sans-serif; font-size: 108%;"> - Some elements in Group IV and some of their compounds - Examples: Diamond, graphite, silicon dioxide

__Types of forces present__ - Strong covalent bonds - (electrostatic) force between the nuclei of the atoms and the shared electrons __Volatility__ - High melting point - High boiling point Reason: Atoms in the lattice are held by strong covalent bonds. A lot of energy is needed to overcome these forces. __Hardness__ - Very hard (if 3 dimensional) __Electrical conductivity__ - Do not conduct electricity in any state Reason: Made up of neutral atoms, thus there are no mobile charged particles to carry charges and conduct electricity Exception: Graphite is the only non-metal that is a good electrical conductor Reason: Each carbon atom is bonded to 3 other in the graphite lattice. The 4th valence electron is delocalized within each layer and act as a mobile charge carrier to conduct electricity. __Solubility in water__ - Not soluble in any type of solvent

<span style="font-family: Arial,Helvetica,sans-serif; font-size: 120%;">** __Metallic Lattice__ ** <span style="font-family: Arial,Helvetica,sans-serif; font-size: 108%;"> - Metals - Examples: Na, Mg, Fe, Cu

__Types of forces present__ - Strong metallic bonds - (electrostatic) force between metal ions and delocalized electrons __Volatility__ - Relatively high melting point - Relatively high boiling point Reasons: Electrostatic forces of attraction between the positive metal ions and sea of delocalized electrons are strong and a lot of heat energy is required to overcome these forces. Exceptions: Group 1 metals and mercury have lower melting point and boiling point. __Hardness (Compare with ionic lattice)__ - Hard - Malleable Reason: The metallic bond is strong but not rigid. When a force is applied, the metal ions slide over one another without disrupting the metallic lattice. __Electrical conductivity__ - All metals conduct electricity in the solid and liquid states. Reason: Metallic lattice consists of positive ions surrounded by a sea of electrons. These delocalized electrons can act as a charge carrier to conduct electricity. __Solubility in solvents__ - Not soluble in any type of solvent - However, many metals react with water (thus, appearing to dissolve)

<span style="-webkit-text-decorations-in-effect: none; color: #96e250; font-family: arial,helvetica,sans-serif; font-size: 16px; line-height: 23px;">__**Polarity and electronegativity**__ - In a covalent molecule, atoms are held together because their nuclei are both attracted to the shared electron pair. - Different atoms attract the bonding electrons unequally. - Atoms with a slight negative charge (-) = has a greater share of bonding electrons - Atoms with slightly positive charge (+) = has lost some of its share of bonding electrons
 * __Polar Bonds__**

- Measure of tendency of an atom to attract a bonding pair of electrons - Aroms with strong electron pulling power = highly electronegative - Difference in electronegativity values can be used to predict how polar a particular bond is - Highest electronegativity = Reactive non-metals - Lowest electronegativity = Highly reactive metals
 * __Electronegativity__**

Polar bonds are thus like covalent bonds with a bit of ionic character in them. The ionic and covalent models are extreme forms of bonding; Polar bonds are somewhere in between them. The greater the difference in electronegativity of the 2 atoms, the higher the polar bond and the higher the ionic character.

__** Diamond and Graphite **__ - Made of only carbon atoms that are packed in macromolecular structures - Allotropes = Different forms of the same elements - Allotropes of carbon = Diamond and Graphite

__Melting point__ 3700 degree Celsius - High melting point Reason: Strong covalent bond between atoms, thus more force is needed to break the bonds __Density__ 3.5g/cm3 __Appearance__ - Colourless, transparent crystals __Hardness__ - Hardest natural substance known - Used as drill tips for drilling equipments and in glass cutters Reason: In diamond, each carbon atoms is joined tetrahedrally to 4 other carbon atoms by strong covalent bonds, forming a giant network (lattice). The highly symmetrical structure and strong C-C bonds make diamond very hard as a substance. __Electrical conductivity__ - Does not conduct electricity Reason: In diamond, each of the carbon's 4 valence electrons are involved in covalent bonding with other carbon atoms. Thus, there are no delocalized electrons to more through the structure to conduct electricity.
 * __Diamond__**

__Melting point__ 3300 degree Celcius __Density__ 2.2g/cm3 __Hardness__ - Soft - Used as a solid lubricant to reduce friction ni engines and also as a pencil lead Reason: Graphite's structure, by contrast, consists of giant flat (2 dimensional) layers of carbon atoms. Althoughm the atoms are covanlently bonded within each layer, the bonds between the layers are weak (intermolecular bonds), and allow the layers to slide over one another. This makes graphite soft and slippery. __Electrical conductivity__ - Conducts electricity Reason: In graphite, each carbon atom is covalently bonded to 3 other carbon atoms. This leaves each carbon atom with 1 valence electron not involved in bonding. This electron becomes delocalized and can move more freely along the layers of carbon atoms, thus conducting electricity.
 * __Graphite__**