Chemical bond defined: A bond is a link caused by an attraction between two atoms within a molecule or compound. There are 2 broad categories of bonds: Ionic Covalent

Ionic Ions are created by an atom’s gain or loss of electrons during the bonding process. Covalent This type of bond results from the sharing of an electron pair between 2 atoms. After ions are formed there is an electrostatic attraction between the two ions resulting in the formation of a bond between them. The sharing can be equal or unequal.

Ionic Bonds These types of bonds are created as a result of the formation of two types of charged particles called ions. The two ion types are: Anions: Negative ions created when neutral, ground-state atoms gain one or more electrons. Cations: Positive ions created when neutral, ground-state atoms lose one or more electrons.

Ionic Bonds These types of bonds can be predicted using three methods: 1. Position in the Periodic Table, (predictor) Atoms from two different families separated by a great distance tend to form an ionic bonds. 2. Metal bonding with a non-metal, (predictor) When a metal bonds with a non-metal the bond will most likely be ionic. 3. Difference in electronegativity, (confirmation) When the electronegativity difference is 1.7 or greater the bond will be ionic.

Examples: Sodium and chlorine Sodium Valence of +1 Electronegativity of.93 Chlorine Valence of -1 Electronegativity of 3.16 What type of bond will result? Ionic, due to an electronegativity difference of 2.23.

Sodium Sodium’s valence is +1, which means when bonding with a nonmetal, its tendency is to lose one electron and form the [Na] +1 ion. Sodium does this is to arrive at an electron configuration which resembles it’s nearest Noble gas neighbor Neon, which has a full octet of valence electrons. Chlorine Chlorine’s valence is –1, which means when bonding with a metal, its tendency is to gain one electron and form the [Cl] -1 ion. Chlorine does this is to arrive at an electron configuration which resembles it’s nearest Noble gas neighbor Argon, which has a full octet of valence electrons.

Sodium chloride This is what occurs during the formation of sodium chloride: Na Cl Since Cl’s electronegativity is so much higher than Na’s, it will want to pull the electron away from sodium, (and sodium will allow it), creating two ions, The cation, [Na] +1, and the anion [Cl] -1 [ ] [] +1 This electrostatic attraction creates a bond resulting in the formation of the compound, Cl Na

Bond Types Ionic Bonds Bonds which result from the Transfer of electrons

Bond Types Ionic Bonds

Please Note: The formula for the ionic compound sodium chloride. is NaCl. It has an overall charge of “0”. [Na] +1 + [Cl] -1 = 0 Now, let’s consider the compound Magnesium Iodide: [Mg] +2 + [I] -1 = +1 In order for the overall charge on this compound to equal 0, 1 additional iodine atom is needed. The formula for this compound would be MgI 2 The subscript 2 indicates 2 Iodine atoms

Bond Types Covalent Bonds which result from the sharing of electrons

Bond Types Covalent Bonds

Bond Types Covalent Bonds can be Polar or Non-polar. Covalent Bonds: electronegativity difference of Polar Covalent Bonds: electronegativity difference of

The difference between a polar (water) and non-polar (ethane) molecule is created by an unequal sharing of electrons within the polar molecule. Non-polar molecules have electrons equally shared within their covalent bonds. Polarity

Bond Types Metallic Bonds which only occur in metals and result from the sharing of valence electrons within the metallic substance.

Bond Types Metallic Bonds

Properties of compounds of different bond types Ionic compounds are formed when atoms with very high electronegativities take 1 or more electrons from atoms with very low electronegativities. * If an element from Family 1A or 2A reacts with an element from families 5A or 6A, an ionic compound will result. Ionic compounds are normally solids at room temperature. Ionic Compounds Covalent Compounds Covalent compounds tend to be larger than ionic compounds and are formed by the reaction of atoms of similar electronegativities. Covalent compounds tend to be gases or room temperature. Covalent compounds containing carbon tend to be solids at room temperature and can be very large chain molecules.

Hydrogen Bonds Hydrogen bonds are weak forces of attraction between 2 neighboring polar molecules within a substance. Example: Water Since the electronegativity of oxygen is 3.5 and hydrogen is 2.2, water molecules are polar-covalent. This electronegativity difference creates a condition called a “dipole”. A dipole is caused by a shift of electrons toward the more electronegative atom. This causes the area around the more electronegative atom to become more negative caused by increased electron density and the area around the less electronegative element becomes more positive due to the shift of electrons away from it.

Hydrogen Bond H: electronegativity 2.2 O: electronegativity 3.5 The electronegativity differences between oxygen and hydrogen cause a shift in electron density toward the oxygen portion of the water molecule. This shift of electrons creates a “dipole”. Since the Oxygen is partially negative (δ - ), and the Hydrogen is partially positive (δ + ) due to the dipole, there is an attraction between the oxygen of one molecule and the hydrogen of an adjacent molecule. This attraction is called a hydrogen bond. Hydrogen Bonds, cont. (δ-)(δ-) (δ + )

Metallic Compounds Metallic Compounds are well known for being solids at room temperature, with 1 exception, mercury. Each metallic atom tries to be associated as close as possible with other metallic atoms within the compound. The electrons within the metallic substance are not associated with just one of the metal atoms but can easily move throughout the entire substance. This characteristic is what makes metals good conductors of electricity.

The concept of “Valence.” Valence is a numerical way to represent the “Oxidation state” of an element. This number represents the tendency of the element to gain or lose electrons when bonding with another element.

“Valence”, cont. Valence numbers, referred to as “oxidation states” are charges assigned to each element in Families 1A-8A, on the periodic table. Transition metals contained in the d-block, as well as all f-block elements have multiple valence #’s (oxidation states) and will behave differently depending upon the conditions of the reaction.

“Valence”, cont. Whenever an element is found in it’s oxidation (ion) state, it is called a “monatomic ion” due to the fact that it has lost or gained at least 1 electron in order to assume its charged atomic structure. Example: Cl o gains 1 e - forming the [Cl] -1 ion.

Lewis Structures Lewis Structures A way of showing how elements bond within a compound using the symbols of the elements and dots or lines to represent the bonds between them. Let’s learn how to draw Lewis structures…

Lewis Structures Lewis structures are diagrams which illustrate the arrangement of all bonding (valence) electrons found around each element within a compound. These diagrams must account for all valence electrons for each atom within the compound and will illustrate how the “octet rule” applies in the formation of chemical bonds

Lewis Structures Xx Valence electrons are represented by placing dots on each of the 4 sides of the element’s symbol as if the symbol was in the center of an imaginary box. 2s 1 2s 2 2p 1 2p 2 2p 3 2p 4 2p 5 2p 6 Individual Elements

Lewis Structures Lewis structures for single elements can have up to 8 electrons (2/side) around them. Br Bromine is found in Family 7A & has 7 valence electrons. It’s Lewis structure would be drawn like this:

Sodium Chlorine Draw the Lewis diagram for a single sodium atom: Na Cl Draw the Lewis diagram for a single chlorine atom: Lewis Structures

When drawing Lewis structures for compounds there can be no more than 8 electrons (octet) placed in pairs around any element within the compound. C H In the compound Methane (CH 4 ), each carbon atom brings 4 valence electrons and each hydrogen brings 1 valence electron during bonding. It’s Lewis structure would contain 8 electrons & will look like this: H H H There can be up to 3 pairs of electrons located on any one side of an element’s symbol, but all valence electrons for each element must be accounted for without adding any additional valence electrons.

Lewis Structures CH Within the Methane molecule (CH 4 ), the carbon atom wants to have a full valence shell containing 8 electrons and each hydrogen wants 2. Any dots, (electrons) that are located in between 2 atoms within the compound, count for both atoms. H H H Bonds always involve electron pairs:

Lewis Structures CH Compounds will always try to arrange themselves in a symmetrical manner. In compounds containing more than one of the same element, the element of higher quantity will usually arrange itself around the the central single atom, especially if the single element is Carbon as shown below. H H H Notice the symmetrical arrangement of atoms in this methane molecule:

Lewis Structures Rules for drawing Lewis structures:  Covalent Compounds must show the Valence electrons for all elements in the compound.  Ionic Compounds only need to show the Valence electrons (full octet) for the element with the highest electronegativity, (the nonmetal).  Covalent compounds can have, 1, 2, or 3 electron pairs between 2 atoms to represent single, double or triple bonds.

Refer to the Handout provide on the Rules for Drawing Lewis Structures for a more detailed description

Lewis Structures Double Bond, (2 electron pairs) C H H C H H Each carbon brings 4 valence electrons & each hydrogen brings one, making a total of 12 valence electrons. Since carbon has a need for 8 valence electrons, the only way to reach the proper configuration is to place 4 electrons between the two carbon atoms creating a double bond. C 2 H 4 : Ethane

Draw the Lewis structure for MgI 2 Mg I I I I I I I2I2 I2I2 Lewis structure for each iodine atom Lewis structure for the Magnesium atom Lewis structure for the Magnesium Iodide Lewis structure must contain 16 valence electrons

Now, you try it! Draw Lewis structures for these simple compounds, ( use different colors or X’s and 0’s to represent each element’s valence electrons ). Lithium Chloride Rubidium Sulfide (-2) Potassium Fluoride Sodium Oxide Li Cl Rb S S LiCl Lewis structure Formula Rb 2 S F F K K KF Na O O Na 2 O Lewis Structures, cont. Carbon dioxide C C O O O O CO 2

Lewis Structure Activity Pick up a handout on the counter Read the introduction, background and the procedure. After completing the reading, get a partner. Pick up a bag of materials and complete the elements found on Data Table I. ** Please complete the list in order and make sure you and your partner understand how to complete each Lewis structure before moving on to the next. Have Mr. A. check #’s 1-10 as you complete them.