Chemical Bonding 2014

What is a chemical bond? In order to bond atoms must interact. Atoms interact to achieve a stable octet. Atoms will lose, gain or share electrons to get an octet. This creates a bond. The type of bond depends on what takes place with the valence electrons.

What is a chemical bond? Definition: A force of attraction caused by a loss, gain or sharing of electrons that holds atoms together.

Energy of Bonds Chemical energy (potential energy) is released when bonds are formed. Chemical energy is added to break these bonds. Forming a bond lowers potential energy. weak bonds – need little energy to break strong bonds – need a lot of energy to break.

Types of Bonds 1. Ionic – formed by metal with non-metal or positive ion and negative ion 2. Covalent – between two non-metals 3. Metallic – force of attraction that hold atoms of metals together

Ionic Bonds – between ions. Transfer of electrons form ions (charged particles) cations and anions attract each other by electrostatic forces. For example: metal forms + ion (oxidation) Na  Na +1 (loses 1 e-) non-metal forms – ion (reduction) Cl  Cl-1 (gains 1 e-) Lewis dot diagrams show how electrons are transferred.

Example: sodium reacts with chlorine Cl Na Has one valence electron, wants to get rid of it so that its valence shell is full Has 7 valence electrons, needs one more to have a full valence shell This reactions makes sodium chloride or table salt. Na+1[ Cl ]-1

Ionic Bonds – Ionization Energy and electronegativity In order for an ionic bond to be formed one atom must have a low ionization energy (low electronegativity) and the other must have a high ionization energy (high electronegativity) p. 2

Reaction to form an ionic compound In your notes, describe the reaction. Are the reactive properties of these elements consistent with what we learned last chapter about the groups of the periodic table?

Ionic Bonds – electronegativity p. 2 Electronegativity values are used to determine bond type. Differences equal to or greater than 1.7 are considered ionic. Example: Look up on Table S Na = 0.9 Cl = 3.2 3.2 – 0.9 = 2.3 since 2.3>1.7 ionic The higher the END the more ionic the bond.

Try These: Determine the electronegativity difference (END) for the following: KCl MgF2 Which is more ionic???

Try These: Answers Determine the electronegativity difference (END) for the following: KCl 3.2 – 0.8 = 2.4 ionic MgF2 4.0 – 1.3 = 2.7 ionic MgF2 is more ionic than KCl

Exceptions to END > 1.7 H = 2.2 END = 2.2 – 0.9 = 1.3 Metal Hydrides (metals bonded to hydrogen) Example: sodium hydride NaH Na = 0.9 H = 2.2 END = 2.2 – 0.9 = 1.3 Considered ionic because electrons are transferred

Properties of Ionic Solids High melting points Do not conduct electricity, in solid phase. Crystalline structure. As aqueous solutions or when heated ions are free to move so only then will they conduct electricity. Examples: NaCl and MgO

Try These: p. 3 Write Lewis dot structures for the following ionic compounds: Barium oxide Magnesium chloride Aluminum oxide

Covalent Bonds – molecular compounds A bond that exist between atoms caused by a sharing of electrons. Usually occurs between two non-metals (neither is willing to lose electrons) Electronegativity difference < 1.7 Example: H2 CO2

Two Hydrogen Atoms The valence shells overlap and the electrons are shared making a more stable molecule.

REVIEW: Diatomic molecules are atoms of the same element that covalently bond to meet the octet rule. The following elements are diatomic: H2, O2, F2 Br2, I2, N2, Cl2 Remember the trick to remember them: HOF BrINCl

Exceptions: Hydrogen fluoride, HF, has an electronegativity difference of 4.0 – 2.2 = 1.8 This appears to be ionic but is really covalent.

Double Covalent Bonds p. 3 Occurs when atoms share two pair of electrons (4 electrons). Examples: CO2 C2H4 O2

Triple Covalent Bond p. 4 Occurs when atoms share three pair of electrons (6 electrons). Examples: N2 HCN

Important Bonding Patterns The following non-metals can bond in various ways: H O N C Halogens 1 bond 2 bonds 3 bonds 4 bonds 1 bond

Dash structures: (structural formulas) Dashed structures only exist for covalently bonded molecules Each shared pair of electrons is represented by a dashed line ( ) Single bond: one dash ( ) Double bond: two dashes ( ) Triple bond: three dashes ( )

Drawing Lewis Dot Diagrams for Moleculars Add up all the valence electrons for an atom in the molecule. The “central atom” is the one that is most electronegative or there will only be one of this atom. Add the electrons until you reach the total remembering that all the atoms must obey the octet rule except for hydrogen which obeys the duet rule. Also remember that one pair of electrons between two atoms is a single bond two is a double and three pairs is a triple bond. These are known as bonded or shared pairs of electrons. Not every electron has to be bonded. We call these the lone or non-bonded pairs because they are “lonely”.

Lewis dot structures Symbol Valence electrons Move electrons to satisfy octet/duet for all atoms involved Try to write dashed structures for each….

Try These: H2O NH3 CCl4

Drawing Lewis structures for polyatomic ions. Figure out the total number of electrons considering charge (add up all valence electrons from each atom and adjust according to the overall charge) Example: SO4 2- OH- PO4 3-

Drawing Lewis structures for polyatomic ions. Figure out the total number of electrons considering charge (add up all valence electrons from each atom and adjust according to the overall charge) Example: SO4 2- 32 electrons total OH- 8 electrons total PO4 3- 32 electrons total

Lewis Dot Structures for polyatomic ions Symbol Valence electrons Add or subtract electrons based on overall charge Move electrons to satisfy octet/duet for all atoms involved Brackets and charge

SO4 2- OH- PO4 3-

Drawing Lewis structures for polyatomic ions. -2 SO4 2- O O S O O OH-1 PO4 3- -1 O H -3 O O P O

Coordinate Covalent Bonds p. 6 When both electrons in a covalent bond are contributed by one of the two atoms or when an electron is donated by the environment. Example: ammonium ion NH4 +1 sulfate SO4 2-

Resonance p. 5 Two alternate Lewis structures are equivalent except for placement of electrons are called resonance forms. Example: To describe ozone O3 properly, the real molecule is an average of resonance structures. Another example: NO3

Metallic Bonding p. 5 Metals have a crystalline lattice structure in which positive ions are set in fixed positions, valence electrons are free to move belonging to the crystal as a whole not to individual atoms. “sea of electrons” Since electrons can move freely, metals are good conductors of heat and electricity.

Metallic bond

Polarity of Bonds (covalent only) Ionic bonds by nature are always polar: END >1.7 A polar covalent bond occurs when two atoms do not share the electrons in a bond equally, resulting in a partial charge on each end of molecule. This unequal sharing is a result of a difference in electronegativity. END of 0.4-1.6: polar covalent bond p. 6

Examples: HF , H2S , PI3 END: 1.8 0.4 0.5 * HF is an exception to ionic/covalent bonding END rule

Non-Polar Covalent Bonds Occurs when two atoms share electrons in a bond equally. This is a result of little or no difference in electronegativity. END of 0.0 –0.3 Examples: diatomics F2 H2 N2 END’s are always zero!!!!!!!!!!!

Practice: Indicate whether the following have polar covalent, non-polar covalent or ionic bonds. CaO NH3 CO2 PCl3 Br2 HF

Practice: Answers Indicate whether the following have polar covalent, non-polar covalent or ionic bonds. CaO ionic NH3 polar covalent CO2 polar covalent PCl3 polar covalent Br2 non-polar covalent HF polar covalent (exception to 1.7)

Polarity of Molecules p. 6 Polar molecule (dipole) : a molecule that has a slightly + end and a slightly – end. In order for a molecule to be polar it must have polar covalent bonds and asymmetric shape . asymmetric shapes: bent, pyramidal, linear without a central atom Examples: HF H2O NBr3

Linear without a central atom: asymmetrical H Cl

Bent: asymmetrical O H H Must remember non-bonding electron pairs: they still have a repulsion force O H H

pyramidal: asymmetrical Must remember non-bonding electron pairs: they still have a repulsion force N H H H

Polar molecules Examples: HF H2O NBr3

Non-polar Molecules p. 7 A molecule that does not have a slightly + and – end. **Non-polar molecule with non-polar bonds Examples: F2 , H2 , Br2 **Non-polar molecule with polar bonds can be non-polar if it has symmetry: the same on both sides or ends symmetric shapes= tetrahedral, linear with a central atom, trigonal planar examples: CH4 , CO2

Diatomics All non-polar bonds in linear symmetrical : non-polar molecules. Examples: H H F F

Tetrahedral: symmetrical H C H

Linear with central atom: symmetrical O = C = O

trigonal planar: symmetrical Cl Al Cl

http://www.youtube.com/watch?v=QwpH0fEwwmo

Intermolecular forces- p. 7 Definition: A force of attraction between two molecules Related to many physical properties

Types of Intermolecular Forces Dipole-dipole Attraction Hydrogen Bonding Van Der Waals Dispersion Forces

Hydrogen Bonding Attraction between the H of one molecule and the F, O, N of another molecule. Strongest intermolecular force Examples: HF H2O NH3

Dipole-dipole Attraction Attraction between polar molecules Examples: HCl H2S

Van Der Waals Dispersion Forces Attraction between non-polar molecules or between monoatomic atoms Only attraction for non-polar substances Examples: F2 Ne

Van Der Waals Dispersion Forces Van Der Waal Forces increase with: Decreasing distance Increasing mass p. 8

Ionic Crystals p. 8 Repeating particles in crystals are three-dimensional and formed by linking ion pairs. Examples: sodium chloride NaCl Properties: High melting/boiling points Hard Do not conduct electricity as solid Very soluble (dissolve easily) Does conduct as aqueous

A Collection of Crystals

Molecular Crystals p. 8 Repeating particles are electrically neutral. Ex: sugar C6H12O6 Properties: Low melting/boiling points Usually below 300 C High to low solubility Poor conductor Made of nonmetals

Covalent Network Solids Repeating patterns are atoms covalently bonded Ex: diamond, graphite, SiC Properties: Very high melting points Very hard, abrasive Insoluble

Allotropes Different molecular forms of an element in the same physical state Examples: diamond and graphite Oxygen (O2) and ozone (O3)

Metallic Solid Repeating particles in crystal are positive with “mobile valence electrons” moving about the crystal Good conductor of heat and electricity Examples: copper, gold, silver

P Cl Cl Cl