Chemical Bonding. Chemical Bonding 20/03/2015. The atomic radius increases from right to left. The atomic radius increases from top to bottom

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1 Chemical Bonding Atomic Radius: This distance from the nucleus to the outermost electron. Chemical Bonding Chemistry 11 Two factors must be taken into consideration in explaining this periodic trend: Increasing nuclear charge. Increasing number of shells. Along a Period (left to right): The atomic number increases (more protons) while the valence electrons remain in the same shell. Due to the increasing nuclear charge (pulling electrons closer to the nucleus): Along a Group (top to bottom): The atomic number continues to increase. However, the shell increases from shell 1 to shell 2 etc. The atomic orbitals for each successive shell gets larger and larger. The result is: The atomic radius increases from right to left The atomic radius increases from top to bottom 1

Things change when you take ions into consideration: Cations: Get smaller because you lose outer electrons and the net positive charge draws in remaining electrons.

2 Things change when you take ions into consideration: Cations: Get smaller because you lose outer electrons and the net positive charge draws in remaining electrons. Anions: You gain electrons and they repel each other. The ionic radius expands out to accommodate the repulsive forces. Ionization Energy The energy needed to remove an electron from the atom. It is a minimum for the alkali metals which have a single electron outside a closed shell. It generally increases across a row on the periodic table with a maximum for the noble gases which have closed (full) shells. For Example: For the Sodium atom: For Chlorine atom: 2

Electron Shielding Imagine a campfire Electronegativity A chemical property that describes the ability of an atom to attract electrons towards itself in a

The same thing happens in the atom: When a new orbital is started, every orbital of lower energy shields these electrons from feeling the full nuclear charge.

7 to 4.0. If an atom has a HIGH electronegativity, it strongly attracts electrons from a neighboring atoms and may completely remove it.

3 Electron Shielding Imagine a campfire Electronegativity A chemical property that describes the ability of an atom to attract electrons towards itself in a covalent bond. The same thing happens in the atom: When a new orbital is started, every orbital of lower energy shields these electrons from feeling the full nuclear charge. First proposed by Linus Pauling in He developed the Pauling Scale. It uses a dimensionless number from 0.7 to 4.0. If an atom has a HIGH electronegativity, it strongly attracts electrons from a neighboring atoms and may completely remove it. Atoms with high electronegativies also STRONGY ATTRACT THEIR OWN VALENCE ELECTRONS. As a result, these valence electrons are difficult to remove and the atom has HIGH ionization energy. Electron Affinity The energy released when 1 mole of gaseous atoms each acquire an electron to form 1 mole of gaseous 1- ions. The above represents the first electron affinity (E ea ) 3

Lewis Theory Elements of his theory: From 1916-1919, Gilbert N. Lewis made several important proposals on bonding which lead to the development of Lewis Bonding Theory.

This leads to the ion formation and IONIC BONDS Sometimes bonding involves sharing electrons between atoms, and leads to COVELENT BONDS.

Usually this is that of a noble gas This arrangement is called an OCTET (s &p, not d) Lewis Symbols Hold the phone A common chemical symbol surrounded by up to 8 dots.

4 Lewis Theory Elements of his theory: From , Gilbert N. Lewis made several important proposals on bonding which lead to the development of Lewis Bonding Theory. Valence electrons play a fundamental role in chemical bonding. Sometimes bonding involves the transfer of one or more electrons from one atom to another. This leads to the ion formation and IONIC BONDS Sometimes bonding involves sharing electrons between atoms, and leads to COVELENT BONDS. Electrons are transferred or shared such that each atom gains a more stable electron configuration. Usually this is that of a noble gas This arrangement is called an OCTET (s &p, not d) Lewis Symbols Hold the phone A common chemical symbol surrounded by up to 8 dots. The symbol represents the nucleus and the electrons of the filled inner shell orbitals. The dots represent the valence electrons. Covalent Bonds Sharing electrons, strong bond Non-metal and non-metal Cannot conduct electricity if dissolved in water (aq) Most are gas or liquid at STP If solid at STP, = crystals (sugar) Form molecules Rules to building a covalent bond. 1. Calculate the # of valence electrons and add them together 2. Subtract 2 electrons for each bond / write one bond between each atom 3. The remaining electrons must go on the outside atoms 4. If there are any electrons left over, they must go on the central atom 5. Each atom must have 8 electrons except hydrogen (2 electrons) 4

5 ELEMENT # Electrons Lewis Dot Diagram Structural formula ELEMENT # Electrons Lewis Dot Diagram Structural formula H 2 O SO 2 ELEMENT # Electrons Lewis Dot Diagram Structural formula ELEMENT # Electrons Lewis Dot Diagram Structural formula NO 3 - *** HCN The Lewis Structures of Covalent Compounds that Violate the Octet Rule Incomplete Octets: 1. Odd number of electrons ex. NO The Lewis Structures of Covalent Compounds that Violate the Octet Rule (cont.) Incomplete Octets: 1. Odd number of electrons 2. Molecules or polyatomic ions that have less than an octet. H, Be (family), and B (family) are exceptions Since they have very low electro negativities, they can only accept one electron for every one the donate Ex. BF 3 5

6 Ex. BF 3 The Lewis Structures of Covalent Compounds that Violate the Octet Rule (cont.) Incomplete Octets: 1. Odd number of electrons 2. Molecules or polyatomic ions that have less than an octet. 3. Hypervalent Compounds Elements in the 3 rd and 4 th periods MAY obtain more than 8 electrons (electrons are placed in low laying d orbitals) The central atom will contain the extra electrons Ex. PCl 5 Ex. PCl 5 Ionic Bond Don t share electrons Association of ions (electrostatic attraction) Strong bond (most of the time it s stronger than covalent) Most compounds are solid at room temp Forms crystals When dissolved in water, they conduct electricity and dissociate. Eg. Made of 1. Metal and non-metal 2. Metal and polyatomic ion 3. Polyatomic ion and non-metal 4. Polyatomic ion and polyatomic ion Polar Covalent Bonds Imagine we have one atom with a somewhat higher electronegativity than the other in a covalent bond: This will cause the electrons to be shared unevenly, such that the shared electrons will spend more time (on average) closer to the atom that has the higher ELECTRONEGATIVITY. 6

Ex. H 2 O The greater the difference in electronegativity in the bonding atoms, the greater the polarity of the bond.

True NON-POLAR COVALENT BONDS form only when diatomic molecules are formed with two identical atoms (ΔE 0.

9) Hydrogen Bonding Polar molecules, such as water molecules, have a weak, partial negative charge at one region of the molecule (the oxygen atom in

7 Ex. H 2 O The greater the difference in electronegativity in the bonding atoms, the greater the polarity of the bond. Atoms with a widely different electronegativity values (ΔE 2.0) tend to form IONIC BONDS. True NON-POLAR COVALENT BONDS form only when diatomic molecules are formed with two identical atoms (ΔE 0.4) Everything else will form a POLAR COVALENT BOND (ΔE ) Hydrogen Bonding Polar molecules, such as water molecules, have a weak, partial negative charge at one region of the molecule (the oxygen atom in the water) and a partial positive charge elsewhere (the hydrogen atoms in water). When water molecules are close together, their positive and negative regions are attracted to the oppositely charged regions of nearby molecules. The force of attraction, shown here as a dotted line, is called a HYDROGEN BOND Each water molecule is hydrogen bonded to four others: HYDROGEN BONDS are WEAK 7

8 London Dispersion Forces The London dispersion force is the weakest intermolecular force Because of the constant motion of the electrons, an atom or molecule can develop a temporary (instantaneous) dipole when its electrons are distributed unsymmetrically about the nucleus It is a temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles. Dispersion forces are present between any two molecules (even polar molecules) when they are almost touching kristine LONDON DISPERSION FORCES are the WEAKEST 8

Periodic Trends. Homework: Lewis Theory. Elements of his theory:

Periodic Trends. Homework: Lewis Theory. Elements of his theory: Periodic Trends There are various trends on the periodic table that need to be understood to explain chemical bonding. These include: Atomic/Ionic Radius Ionization Energy Electronegativity Electron Affinity