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CHEMICAL BONDING

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Elements create bonds in order to obtain stable gas configuration through the gain, loss, or sharing of electrons.

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Covalent bond – Electrostatic forces of attraction between positively charged nuclei and shared pair of electrons.

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Metallic bond – Electrostatic forces of attraction between positive metal ions and a sea of delocalized electrons.

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Ionic bond – The electrostatic attraction between oppositely charged ions. Ionic bonds are formed by elements with an electro-negativity difference of 1.8 units or larger.

Ionic Bonding

  • Non-directional

  • Giant Ionic Lattice Structure

 

 

 

Valency – Combining power of an element

Coordination number – Expresses the number of ions that surround a given ion in a compound.

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e.g. 

Ammonium chloride NH4Cl (Both non-metals make this compound. It features two types of bonding – ionic between the ions and covalent bonding between the atoms that make up NH4.)

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HOW ARE IONIC BONDS FORMED? 

Complete transfer of electrons from a metallic atom to a non-metallic atom, to form oppositely charged ions which attract each other.

STRENGTH OF AN IONIC BOND 

1. Charge of the ion (larger - stronger) 

2. Size of the ionic radius (smaller - stronger)

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MAGNESIUM CHLORIDE

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LITHIUM NITRIDE

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LEWIS STRUCTURE 

Shows all valence electrons in a covalently bonded species. 

 

RULES 

1. Add total valence electrons 

2. Draw skeletal structure 

3. Use a line between each element to symbolize an electron pair 

4. If not enough electrons, use double/triple bonds 

5. Ions must have square brackets around them with a charge.  

SODIUM OXIDE

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Covalent Bonding

In covalent bonds, strength increases if: 

  • bond length decreases 

  • shared electrons increases  

(Orbitals must be completely empty or half filled to take place in bonding)

BEN - Breaking bonds is endothermic 

MEX - Making bonds is exothermic

tendency to attract a shared pair of electrons to itself. 

BOND STRENGTH - How much energy is required to break the bond. 

triple bond > double bond > single bond

SIGMA - Head on overlapping of orbitals

PI - Sideways overlapping 

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*When atomic orbitals overlap, molecular orbitals are formed. 

POLAR BOND - When there is an electro-negativity difference. 

 (i) electronegativity difference 

 (ii) dipole moment (charge & distance)  

is 0 when element is the same. 

MOST ELECTRO-NEGATIVE: N, O, F 

If molecule is symmetrical, it is not polar! 

Electronegativity difference! 

  • Difference of < 1.8 shows bond considered covalent 

  • 0.0 - 0.4 considered non polar 

  • 0.4 - 1.8 considered polar 

  • 1.8 < considered ionic 

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COORDINATE COVALENT BONDS - Atoms must have at least one lone pair. Electrons come from only one atom. 

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RESONANCE STRUCTURES

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OCTET RULE - Refers to the tendency of atoms to gain a valence shell with a total of 8 electrons. 

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Left over electrons are shifted to the central atom (can have more than 8). 

Terminal & Central Placements

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COORDINATE BONDS

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carbon monoxide 

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dimer of aluminum chloride

ammonium cation

VSEPR Theory (Valence Shell Electron Pair Repulsion Theory) 

  • The pair of valence electrons are arranged as far apart as possible from each other. 

  • Focus on central atom and see how many bond and lone pairs there are. 

  • 1 double/triple bond will be considered 1 bond. 

 

1. Electron Domain Geometry - How many bond and lone pairs there are. 

2. Molecular Geometry - Overall shape of the molecule 

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Giant Molecular Structures 

Allotropes - Different forms of an element in the same physical state. ​

                     Different bonding within the structure gives rise to different physical properties. 

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GRAPHITE 

  • Each atom is sp2 hybridized, covalently bonded to 3 others 

  • Hexagons in parallel layers with bond angles of 120 degrees 

  • Layers held together by Van der Waal's forces of attraction, slide over each other 

  • Density: 2.26g/cm3 

  • Contains one non-bonded, delocalized electron per atom 

  • Bad heat conductor 

  • Very high melting point, most stable allotrope

  • Non-lustrous, grey solid 

  • Lubricant, pencils  

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DIAMOND 

  • Each atom is sp3 hybridized, covalently bonded to 4 others tetrahedrally in regular repeating patterns with 109.5 bond angles 

  • Hardest known natural substance 

  • Density: 3.51 g/cm3 

  • All electrons bonded, thus does not conduct electricity 

  • Very good head conductor

  • Very high melting point 

  • Brittle 

  • Lustrous crystal 

  • Jewelry; tools

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GRAPHENE 

  • Each C atom is sp2 hybridized and covalently bonded to 3 other carbons, forming hexagons with bond angle 120. 

  • Two dimensional single later (honeycomb structure) 

  • Density: 1.56 g/cm3 

  • contains one non-bonded, delocalized electron per atom 

  • Conducts electricity 

  • Excellent heat conductor 

  • Very high melting point, stronger than steel 

  • Thinnest material to exist; almost completely transparent 

  • Touch screens; performance electronics 

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FULLERENE (C60)

  • Each C atom is sp2 hybridized, covalently bonded in a sphere of 60 carbon atoms, consisting of 12 pentagons and 20 hexagons 

  • Closed spherical cage in which each C atom is bonded to 3 others 

  • Density: 1.726 g/cm3 

  • Easily accepts electrons to form negative ions; semi-conductor at normal temp and pressure 

  • Very low heat conductivity 

  • Low melting point 

  • Yellow crystaline solid - soluble in benzene 

  • Nanotubes; catalysts, and lubricants 

Metallic Bonding

The electrostatic forces of attraction between a lattice of positive metal ions and a sea of delocalized electrons. 

STRENGTH OF A METALLIC BOND 

  1. No. of delocalized electrons (more - stronger) 

  2. Size of cation (smaller - stronger) 

  3. Charge of cation (larger - stronger) 

PROPERTIES OF METALS 

- Good electricity conductors 

- Good heat conductors 

- Malleable 

- Ductile 

- High melting point 

- Shiny and lustrous (delocalized electrons reflect light!) 

ALLOYS 

1) Steel - Iron + carbon + XYZ 

    High tensile strength, tends to corrode 

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2) Stainless steel - Iron + nickel/chromium 

    High strength, corrosion resistant 

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3) Brass - Copper + Zinc 

    Plumbing 

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4) Bronze - Copper + Tin 

    Coins + Tools 

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5) Pewter - Tin + Antimony + Copper 

    Decorative objects 

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6) Sterling silver - Silver + copper 

    Jewelry + art objects  

Van Der Waal's Forces (Intermolecular forces)

LONDON FORCES 

  • Exists in all molecules 

  • Temporary dipole moment is created due to constant shifting of electron clouds 

  • This instantaneous dipole has an influence on adjacent molecules 

  • Causes the positive nuclei to be attracted to the electron cloud of another molecule, and results in temporary dipoles to be created and molecules to become unsymmetrical 

What affects magnitude? 

- Number of electrons 

- Size (volume) of electron cloud 

- Shape of molecules 

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DIPOLE-DIPOLE FORCES

  • Exists in all polar molecules with a permanent dipole moment 

  • e.g. HF, ICl, HCl, CH3, CHO 

  • Attraction between the positive end of one permanent dipole and the negative end of another permanent dipole on an adjacent molecule. 

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HYDROGEN BONDING 

  • Between most electronegative elements O, N, F and H (e.g. O-H, N-H, F-H) 

  • Examples: ammonia (NH3), hydrogen fluoride ions (HF), water (H2O), (CH3)2O

  • Strongest type of intermolecular forces of attraction 

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DIPOLE-INDUCED DIPOLE 

  • A weak attraction when a polar molecule induces a dipole in an atom or in a non-polar molecule by distrubing thr arrangement of electrons in non-polar species. 

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STRENGTH OF INTERMOLECULAR FORCES! 

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HYDROGEN BONDS > DIPOLE-DIPOLE > DIPOLE-INDUCED DIPOLE > LONDON FORCES 

Ionic Bonding

Covalent Bonding

Covalent Structures

Intermolecular Forces

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